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Gen Chem 1 Flashcards

1
Q

What is the atomic number (Z)

A

the number of protons an atom has

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2
Q

What is the mass number (A)

A

the combined number of protons and neutrons

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3
Q

When the convention showing the atom, it’s atomic number and mass number how is it arranged

A

the mass number on top of the atomic number. both next to the atom name

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4
Q

What are valence electrons

A

electrons farthest away from the nucelus

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5
Q

what happens to valence electrons as they get farther away from the nucleus

A

the farther away they are the less they are attracted to the nucleus, and the more they can interact with other atoms.

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6
Q

What is an ion

A

a positive or negative atom due to the loss or gain of electrons

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7
Q

how many, n’s, p’s, and e-‘s in a Nickel-58 atom

A

Ni = 28 protons
58-28 = 30 neutrons
since it is neutral it has 28 electrons too

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8
Q

How many n’s p’s, and e-‘s does a +2 Nickel-60 atom have

A

Ni = 28 protons
60-28 = 32 Neutrons
Since it is plus 2 it has lost 2 electrons = 26 e-‘s

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9
Q

How much is one amu in grams

A

1.66 x 10^-24

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10
Q

How many amu does carbon have

A

12

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11
Q

what is the atomic weight

A

the weight in grams of one mole of a given element (g/mol)

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12
Q

What is avogadro’s number

A

the number of particles in a mole

6.022 x 10^23

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13
Q

What are isotopes

A

when atoms (having the same number of protons) have different numbers of neutrons

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14
Q

how does isotopicity affect reactivity

A

it doesn’t change a lot because they still have the same number of protons and e-‘s

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15
Q

How is the atomic weight of the elements on the periodic table found

A

you take the mass number of each of the isotopes multiplied by the frequency of that isotope, and add all of those together

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16
Q

Q has
isotope A 60% at 40 amu
isotope B 25% at 44 amu
isotope C 15% at 41 amu

A

40 x .6 = 24
44 x .25 = 11
41 x .15 = 6.15

24 + 11 + 6.15 = 41.15 g/mol

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17
Q

What is plancks constant

A

6.626 x 10^-34 J*s

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18
Q

what is the equation for the energy of a quantum

A 
E = hf
h = plancks constant
f = frequency
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19
Q

What is the equation for the angular momentum of an electron

A

angular momentum = nh/2pie
n = quantum number
h = plancks constant

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20
Q

What is bohrs model of the Hydrogen atom

A

a single proton with an electron circling it

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21
Q

Since angular momentum of an electron has thee different constants in the equation, what determines angular momentum of an electron

A

the quantum number

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22
Q

What is the equation for the energy of an electron

A 
E = -Rh/n^2
Rh = 2.18 X 10 ^-18 j/electron
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23
Q

What is Rh

A

the rydberg constant = 2.18 x 10^-18 j/electron

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24
Q

What is the only factor that affects the energy of an electron

A

the quantum number

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25
Q

What will the energy of an electron be in any of it’s quantized states

A

negatve

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26
Q

The energy of an electron is related to

A

it’s orbital radius the larger the radius the more energy it has. the smaller the radius the less energy it has

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27
Q

what is the smallest orbit an electron can have

A

n = 1

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28
Q

What is the ground state level for an electron

A

when the electron is at it’s lowest energy state

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29
Q

How is atomic emission specrta gathered

A

you take an atom in its ground state, then heat it up or add energy to it and the electron is excited, or moved to a higher energy state. It will quickly move back to it’s ground state but in doing so it will release energy in the form of photons.

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30
Q

What is the equation for the energy of the photons released when doing atomic emission spectra

A 
E = hc/Lambda
H = plancks constant
c = speed of light
Lambda = wavelength
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31
Q

What is the speed of light

A

3 x 10^8 m/s

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32
Q

How are atomic emission spectra used to identify an element

A

each electron in a element when returning to ground state releases a specific wavelength of light, each element has its own specific electrons so each element has it’s own unique atomic emission spectrum

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33
Q

Can the atomic emission spectrum of stars be done

A

ye[p

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34
Q

What is the balmer series of hydrogen emission lines

A

transitions from n > 2 to n = 2

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35
Q

how many wavelengths are in a balmer series of hydrogen emission lines

A

4 in the visible region

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36
Q

what is the lyman series of hydrogen emission lines

A

transitions fron n > 1 to n = 1

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37
Q

what causes an electron to jump up to a higher energy state

A

it absorbs energy

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38
Q

what is an absorption spectrum

A

it is the spectrum showing which wavelengths of energy are absorbed by a specific element

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39
Q

When can absorption spectrum be used

A

in the identification of elements present in a gas phase sample

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40
Q

Why does Bohr’s model not work for atoms with more than one e-

A

because it doesn’t take the repulsion between each electron into account

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41
Q

What is the main difference between Bohr’s model and modern quantum mechanics

A

that electrons don’t go around the nucleus in a circle, but they move around in orbitals

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42
Q

what does an electron orbital show

A

it represents the probability of finding an electron in a given region

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43
Q

what is the Heisenberg uncertainty principle

A

states that it is impossible to determine with perfect accuracy the momentum and position of an electron simultaneously

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44
Q

What are the four quantum numbers

A

n, ℓ, mℓ, ms

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45
Q

What is the pauli exclusion principle

A

no two electrons in a given atom can posses the same set of four quantum numbers.

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46
Q

What does the position and energy of an electron described by its quantum numbers give you

A

its energy state

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47
Q

What is the quantum number n and what does it indicate

A

the principal quantum number

it indicates the distance of the electrons from the nucleus

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48
Q

What is the quantum number ℓ and what does it indicate

A

the azimuthal quantum number (angular momentum quantum number)
it refers to the subshells of the valence electrons

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49
Q 
what do the ℓ quantum numbers
0
1
2
3
indicate
A

0 means that the electrons are in the S subshell
1 means the p subshell
2 = d subshell
3 = f subshell

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50
Q

With using ℓ, how many electrons can fill particular subshells

A

4f + 2

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51
Q

the greater the number of ℓ, the

A

greater the energy of the subshell

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52
Q

what is the quantum number mℓ, and what does it indicate

A

it is the magnetic quantum number

it indicates which orbital

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53
Q

how many mℓ values can there be in terms of ℓ

A

2ℓ + 1

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54
Q

What is the ms quantum number and what does it indicate

A

it is the spin quantum number

and it is just the last way to identify electrons (both electrons in one orbital must have opposite spins

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55
Q

What are electrons with parallel spins

A

electrons in different orbitals with the same Ms values

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56
Q

What are paired electrons

A

electrons in the same orbital with opposite spins

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57
Q

In the electron configuration what does the
1st number
letter
2nd number indicate

A
  1. the principal energy level
  2. the subshell
  3. the number of electrons in that subshell
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58
Q

how do you know which subshells will fill first

A

those with a lower n + ℓ value, if they are tied, then you go with the one with the lower n value

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59
Q

In what order are subshells filled

A 
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
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60
Q

What is hund’s rule

A

all orbitals of the same energy level fill halfway before any will completely fill

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61
Q

What are some exceptions to the orbital filling rules

A

Chromium
- expected = 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4
- actual = 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
This is so that the 3d orbitals can be half filled
Copper
- does the same thing except 5 electrons later, one e- is promoted to the 3d orbital from the 4s obital to give it 10 e-‘s

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62
Q

what are paramagnetic materials

A

materials with unpaired electrons, that are weakly attracted to a magnetic field

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63
Q

What are diamagnetic materials

A

materials with no unpaired electrons that are slightly repelled by a magnetic field

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64
Q

What are valence electrons

A

all of the electrons in it’s outer energy shell

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65
Q

What is the periodic law

A

that the chemical properties of the elements are largely dependent on their atomic number

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66
Q

What are the periods of the p table

A

rows

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67
Q

what are the groups on the p table

A

columns

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68
Q

what does the roman numerals on the periodic table groups indicate

A

the number of valence electrons

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69
Q

What does the A or B indicate after the roman numeral

A 
A = representative elements (either s or p as outermost orbitals)
B = nonrepresentative elements ( have d as their outermost orbitals)
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70
Q

What are the two kinds of non-representative elements

A

transition elements

lanthanide and actinide serie

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71
Q

what are transition elements

A

those with partially filled d orbitals

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72
Q

what are lanthanide and actinide elements

A

those with partially filled f orbitals

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73
Q

what do all elements strive for

A

8 valence electrons

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74
Q

how does atom size radii across the p table

A

as you go from left to right they get smaller

as you go from bottom to top they get smaller

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75
Q

how do they find atomic radii

A

it is one half of the distance between to atoms of that element that are just touching each other

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76
Q

why does atomic radii get smaller towards the top right of the p table

A

from left to right because the more protons and electrons you have the more that they are pulled together
from bottom to top because filled orbitals block the attraction for the valence electrons and the nucleus.

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77
Q

what is the ionization energy

A

the amount of energy required to remove an electron from an atom

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78
Q

how does ionization energy change on the p table

A

the same way that atomic radii does. it is the largest in the top right because the electrons are the most tightly held together

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79
Q

what is first and second ionization energy and how do they compar

A

first ionization energy describes the energy required to remove the first electron, second “ “ second electron
second ionization energy is always greater than the first ionization energy

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80
Q

What is electron affinity

A

the energy change that occurs when an electron is added to a gaseous atom, (it represents the ease with which an atom can accept an electron)

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81
Q

how does atomic size affect electronegativity

A

the smaller the atom, the more it will pull on the electron

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82
Q

how does electronegativity change across the p table

A

it increases towards the top left

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83
Q

what is the effective nuclear charge (Zeff)

A

the attractive pull of the nucleus (increases from left to right)(more protons)

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84
Q

What is electronegativity

A

the measure of attraction an atom has for electrons in a chemical bond.

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85
Q

what is special about cesium

A

-largest
- most metallic
least electronegative of all naturally occurring elements
- smallest ionization energy
- least exothermic electron affinity

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86
Q

What is another way to think about electronegativity

A

nuclear positivity

the stronger the nuclear pull the more electronegative it will be

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87
Q

What is the most electronegative atom

A

fluorine

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88
Q

What are the three types of elements

A

metals
non-metals
metalloids (semimetals)

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89
Q

where are metals, non-metals, and mettaloids located on the p tabel

A

metals left of the ziggy line
non-metals on the right of the ziggy line
mettaloids form the ziggy line

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90
Q

What are the characteristics of metals

A
  • shiny solids at room temperature (except mercury)
  • high melting points
  • high densitties
  • malleable
  • ductile
  • low ionization energies
  • low electronegativity
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91
Q

What is malleability

A

the ability of a metal to be hammered into shapes

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92
Q

What is ductility

A

the ability to be drawn into wire

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93
Q

What gives rise to the characteristic properties of metals

A

the few electrons in the valence shells can easily be removed

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94
Q

What makes metals good conductors of heat and electricity

A

that their electrons can easily be moved.

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95
Q

What are some of the characteristics of non-metals

A
  • brittle in solid state
  • little or no luster
  • high ionization energies and electronegativities
  • gain electrons easily
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96
Q

What are some of the characteristics of metalloids

A
  • posses characteristics of both metals and non-metals
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97
Q

What elements are metalloids

A 
Boron
Silicon
Germanium
Arsenic
Antimony
Tellurium
(polonium)
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98
Q

What are alkali metals

A

elements in group IA

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99
Q

What are the characteristics of alkali metals

A
  • one loosely bound e-
  • highly reactive due to low ionization energies
  • react highly with halogens
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100
Q

What are alkaline earths

A

elements in group IIA

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101
Q

What are some of the characteristics of alkaline earths

A
  • large atomic radii
  • fairly loosely bound electrons 2
  • low electronegativity, ionization energies, and electron affinities
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102
Q

What are halogens

A

elements in group VIIA

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103
Q

What are the characteristics of halogens

A
  • 7 valence electrons

- high electronegativities

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104
Q

What are the physical states of halogens at room temperature

A

F and Cl are gasses
Br is a liquid
I is a solid

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105
Q

What are the noble gases

A

the elements of group VIIIA

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106
Q

what are the charateristics of the noble gasses

A
  • quite unreactive
  • high ionization energies
  • low boiling points
  • all gasses at room temp
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107
Q

What are the transition elements

A

elements in group IB to VIIIB

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108
Q

What are the charateristics of transition elements

A
  • very hard
  • high melting points
  • malleable
  • highly electrically conductive
  • low ionization energies
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109
Q

What are the positively charged transition metals called

A

oxidation states

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110
Q

What causes transition metals to be malleable and conductive

A

they have loosely held d-orbital electrons

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111
Q

What is it called when transition metals in their oxidation states complex with water molecules, or nonmetals

A 
hydration complex (with water)
complexes without
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112
Q

What happens with complexes

A

they absorb a wide variety of frequencies of lights. which leads to interesting colors of liquids.

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113
Q

What interacts to form bonds between molecules

A

the valence electrons of molecules

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114
Q

what is the difference between intramolecular forces and intermolecular forces

A 
intra = the forces that actually hold the two atoms together
inter = other forces within the molecule that are between atoms that aren't bonded
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115
Q

What is the octet rule

A

that atoms bond to other atoms until they have 8 electrons in their outermost electron shell

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116
Q

what are exceptions to the octet rule

A 
Hydrogen  (2 or 4)
Lithium and beryllium (2 or 4)
Boron (6)
elements beyond the second row
phosphorus and sulfur (8+)
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117
Q

What is ionic bonding

A

when electrons from an atom with a small ionization energy are transferred to an atom with a high electron affinity.

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118
Q

what kinds of forces hold ionic bonds together

A

electrostatic

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119
Q

what is covalent bonding

A

When an electron pair is shared between two atoms.

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120
Q

What is it called when the bonds are partially covalent and partially ionic

A

polar covalent bonds

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121
Q

in which kind of bond are electrons completely donated, what kind are they just shared

A

donated in ionic

shared in covalent

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122
Q

how large does the difference in electronegativity need to be for it to be a strong enough ionic bond to create cations and anions

A

at least 1.7 difference or larger

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123
Q

what are the characteristics of ionic compounds

A
  • high boiling and melting points
  • conduct electricity in the liquid and aqueous states (not in solid states)
  • form crystal lattices that increase attractive forces and decrease repulsive forces
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124
Q

What kind of bond is formed when elements of similar electronegativities bind

A

covalent bond

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125
Q

which is requires more energy. the creation of two ions, or the energy released when an ionic bond is broken

A

the energy required to create ionic bonds

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126
Q

what are the characteristics of covalent compounds

A
  • weak intermolecular forces
  • low melting solids
  • do not conduct electricity in liquid or aqueous states
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127
Q

what is bond order

A

it tells you how many electron pairs are shared in a bond
single bond = bond order of 1
double bond = bond order of 2
…

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128
Q

What are the two features of a covalent bond

A

bond length

bond order

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129
Q

how are bond order and bond length related

A

higher bond orders = shorter bond lengths

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130
Q

What is bond energy

A

the amount of energy required to seperate two bonded items

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131
Q

how is bond order and bond energy related

A

the higher the order the higher the bond energy

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132
Q

what is a formal charge

A

when more or less electrons are attributed to an atom in a molecule than would be alone

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133
Q

What are the steps of drawing lewis structures

A
  1. Draw the skeletal structure of the molecule
  2. count all of the valence e-‘s (sum = total number of electrons)
  3. draw single bonds
  4. complete all of the octets (outside atoms first)
  5. put extra electrons on the central atom
  6. double or triple bonds if needed
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134
Q

How do you calculate the formal charge of an atom

A
  1. take the number of valence electrons it normally has
  2. subtract half of all the bonding electrons
  3. subtract all of the non-bonding electrons
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135
Q

What is a resonance structure

A

when two or more non-identical lewis structures can be drawn

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136
Q

how do we know which resonance structure is the actual molecule

A

it’s not one or the other but a hybrid of both, it spends more time, or is more like the more stable of the two resonance structures

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137
Q

What kind of arrow represents resonance structures

A

one double headed arrow

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138
Q

Are lewis structures with formal charges preferred over lewis structures without formal charges

A

no structures without formal charges are more preferred

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139
Q

What are some exceptions to the octet rule

A

atoms found in or below the third period can have more than 8 valence electrons because the D orbital can hold some

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140
Q

What are polar covalent bonds

A

bonds formed between atoms with small electronegativity (.4-1.7)
the electron pair is held more closely to the more electronegative atom
and thus it has a partial negative charge and the other has a partial positive charge

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141
Q

what is a dipole moment

A

the product of charge magnitude and the distance between the partial charges in a polar covalent bond

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142
Q

What are coordinate covalent bonds

A

when the bonding electron pair comes from only one of the two atoms

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143
Q

where are coordinate covalent bonds normally found

A

in lewis acid-base complexes

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144
Q

what is a lewis acid

what is a lewis base

A

lewis base = electron pair donor

lewis acid = electron pair acceptor

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145
Q

2 regions of electron density

shape and degree

A

linear

180 degrees

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146
Q

3 regions of electron density

A

trigonal planar

120 degrees

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147
Q

4 regions of electron density

A

tetrahedral 109.5 degrees

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148
Q

5 regions of electron density

A

trigonal bipyrimidal

90, 120, 180 degrees

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149
Q

6 regions of electron density

A

octahedral

90, 180 degrees

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150
Q

What does an ℓ=0 orbital look like

A

it is an s-orbital

a perfect sphere

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151
Q

What does a ℓ=1 orbital look like

A

that is a p-orbital

there are three along the x,y,and z axis. they are two baloons tied together shape (dumbell)

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152
Q

What does a ℓ=2 orbital look like

A

That is a d orbital

you don’t need to know

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153
Q

What is a molecular orbital

A

when two atomic orbitals interact during a bond

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154
Q

What is the molecular orbital of a single bond

A

both of the orbitals overlap head to head and form sigma bond

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155
Q

What is the molecular orbital of a double bond

A

the orbitals are parallel and it is called a pie bond

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156
Q

How do you get a bonding orbital

A

when the two bonding atomic orbitals are the same sign, they bond

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157
Q

how do you get a non-bonding orbital

A

when the two bonding orbitals are opposite signs, they don’t bond

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158
Q

What are intermolecular forces

A

forces between molecule

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159
Q

What are the kinds of intermolecular forces

A

Dipole-dipole interactions
hydrogen bonding
dispersion forces

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160
Q

What are dipole dipole interactions

A

when the polar (+) end of one molecule orients itself near the polar ( - ) end of another molecule

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161
Q

when do dipole dipole interactions occur

A

in solid and liquid phases

gasses the molecules are too spread out

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162
Q

what can dipole-dipole interactions do for a solid or molecule

A

they can increase the melting and boiling points compared to non-polar substances of the same weight

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163
Q

What is hydrogen bonding

A

when a H bonded to an electronegative atom (F,O,N) and it’s partial positive charge interacts with the partial negative charge on other molecules

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164
Q

what does hydrogen bonding do to the melting point and boiling points of molecules

A

raises them significantly

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165
Q

what kinds of molecules have greater dispersion forces

A

large atomed molecules

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166
Q

what determines if molecules are miscible or immiscible

A

the intermolecular forces

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167
Q

what is a compound

A

a pure substance that is composed of two or more elements in a fixed portion

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168
Q

What is the ionic substitute for molecular weight

A

formula weight

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169
Q

What is the Gram Equivalent weight

A

the molar mass (how many grams per mole a substance is)

divided by the number of (hydrogens) that the substance can supply.

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170
Q

how do you find equivalents

A

weight of the compound / Gram equivalent weight

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171
Q

what good can equivalents of weight do

A

it can help you know how much acid you will need to neutralize a base

172
Q

What is the law of constant composition

A

states that any sample of a given compound will contain the same elements in the same ratio

173
Q

What is the empirical formula

A

the simplest whole number ratio of the elements in a compound

174
Q

what is the molecular formula

A

gives the exact number of atoms of each element in a compound

175
Q

What is the empirical then molecular formula of benzene

A 
empirical = CH
molecular = C6H6
176
Q

What is the percent composition

A

the weight percent of a given element in a compound

177
Q

How do you find percent composition

A

get the mass on the substance in the formula / the formula weight of the compound

178
Q

what are combination reactions

A

two or more reactants form one product

179
Q

what are decomposition reactions

A

one reactant breaks down into two or more products

180
Q

What are single displacement reactions

A

Redox reactions

when one atom or ion moves from one compound to the other

181
Q

What is a double displacement reaction

A

a metathesis reaction

when elements of two compounds react to form two different compounds

182
Q

what is a spectator ion

A

an ion in an ionic equation that doesn’t change from reactants to products

183
Q

what is a net ionic equation

A

an ionic equation in which the spectator ions are left out

184
Q

What are neutralization reactions

A

a double displacement reaction in which water and a salt are created and an acid and base are neutralized

185
Q

What is the mechanism of a reaction

A

the actual steps through which the reaction occurs

186
Q

What is the rate determining step

A

the slowest step of the reactoin

187
Q

what are the two ways to determine the rate of a reactoin

A

the dissapearance of reactants/time

the formation of products/time

188
Q

how do you distinguish between rates by loss of reactants vs. formation of products

A

rate by loss of reactants will be negative

189
Q

what are the units of reaction rates

A

moles/L*s

M/s

190
Q

how do you find the exponents in the rate law equation

A

they are equal to the coefficients of the substances IN THE RATE DETERMINING STEP

191
Q

how do you find the reaction order for a reaction

A

add the exponents of the rate law

192
Q

What is a zero order reaction

A

one in which the rate is independent of the concentrations of the reactans

193
Q

what is a first order reaction

A

the rate is proportional to the concentration of one reactant

194
Q

what is the radioactive decay equation

A 
At = A0 e^-kt
At = concentration at time t
A0 = initial concentration
k = rate constant
t = time
195
Q

what is the half life equation

A

t1/2 = .693/k

196
Q

What is a second order reaction

A

the rate is proportional to the product of the concentration of two reactants, or to the square of one reactant.

197
Q

What are the units for a zero oder reaction

A

Ms-1

198
Q

What are the units for a first order reaction

A

s-1

199
Q

What are the units for a second order reaction

A

M-1s-1

200
Q

What are higher order reactions

A

reactions that have a reaction order greater than 2

201
Q

what are mixed order reactions

A

reactions with a fractional order (X^1/3)

202
Q

What is the collision theory of chemical mechanics

A
  • the rate of a reaction is proportional to the number of collisions per second between reacting molecules
203
Q

What are effective collisions

A

collisions that result in formation of a product

204
Q

what has to happen for a collision to be effective

A

the molecles must be moving fast enough and be properly oriented

205
Q

What is the activation energy

A

the minimum amount of energy needed to break old bonds and form new ones

206
Q

how do you get the rate of reactions with total number of collisions per second (Z) and the fraction of collisions that are effective (f)

A

rate = fZ

207
Q

What is a transition state

A

when the old bonds are beginning to break and the new bonds are beginning to form. the transition state then degrades into the product

208
Q

What is the energy level of the transition state like

A

it is higher than either the products or the reactants (those use energy from a collision to get to that energy)

209
Q

what is the difference between an intermediate and a transition state

A

transition states have a very very short finite lifetime

210
Q

what does ΔH refer to

A

enthalpy change

this shows the difference between the potential energy of the reactants and the potential energy of the products

211
Q

What does a - or + ΔH mean

A

– means that it is exothermic (heat is given off)

+ means that it is endothermic (heat is absorbed)

212
Q

What two things can increase the reaction rate

A

more effective collisions

stabilization of the activated complex (transition state)

213
Q

What four things affect reaction rates

A
  1. concentration
  2. temperature
  3. medium
  4. catalysts
214
Q

When have we reached equilibrium

A

when there is no net change of the concentrations of products and reactants during a reversable reaction

215
Q

What is the law of mass action

A

at equilibrium the forward and reverse reaction rates will be equal

[C]c[D]d / [A]a[B]b = Kc

216
Q

What is Kc

A

the equilibrium constant (except in gasses it is Kp)

217
Q

What is the reactant quotient Q indicate

A

Q is a measure of the degree to which a reaction has gone to completion

218
Q

what is Q equal to

A

[C]c[D]d / [A]a[B]b

219
Q

What are the characteristics of the equilibrium constant Keq

A
  • pure solids and liquids don’t count
  • if Keq is very large when compared to 1 it is mostly products left
  • if Keq is very small compared to 1 is it mostly reactants left
  • if Keq is close to 1 they they are at equilibium, and the concentrations of products and reactants will be very similar
220
Q

What happens to reactions as you increase the pressure

A

the reaction will go to whatever favors less moles, and decrease volume

221
Q

What is an isolated system

A

it cannot exchange energy or matter with the surroundings

222
Q

what is a closed system

A

energy, but not matter can be exchanged with the surroundings

223
Q

what is an open system

A

energy and matter can be exchanged with the surroundings

224
Q

what is an isothermal process

A

the process occurs while the temperature remains constant

225
Q

What is an adiabatic process

A

when no heat exchange occurs

226
Q

what is an isobaric process

A

then the process occurs while the pressure remains constant

227
Q

What is negative and positive heat

A

positive heat is heat absorbed by the system

negative heat is heat lost by the system

228
Q

What are endothermic reactions

A

those that absorb heat

229
Q

what are exothermic reactions

A

those that lose heat

230
Q

how many joules are equal to one calorie

A

4.184 J

231
Q

what is the equation for finding the heat (q) lost or gained by a system

A 
q = mcΔT
q = heat
m = mass
c = specific heat 
ΔT = temperature change
232
Q

in constant volume calorimetry how do you find the q or the reaction

A

q rxn = - (q water + q steel)

233
Q

What does the little circle after enthalpy, entropy, and free energy mean

A

that it is done at standard conditions (25 degrees celcius, 1 atm)

234
Q

What is enthalpy (H)

A

it accounts for the heat absorbed or lost by the system

235
Q

How do you find the enthalpy change

A

ΔH rxn = H products - H reactants

236
Q

what does a +ΔH represent

A

a endothermic reaction (absorbs energy)

237
Q

What is Hess’s law

A

the enthalpies of reactions are additive

238
Q

What is the bond dissociation energy

A

the amount of energy needed to break a particular type of bond in one mole of gaseous molecules

239
Q

how do you use bond energy to calculate ΔHrxn

A

ΔHrxn = (ΔH bonds broken) + (ΔH bonds formed)

240
Q

What is the bond dissociation energy

A

the amount of energy needed to break a particular type of bond in one mole of gaseous molecules

241
Q

how do you use bond energy to calculate ΔHrxn

A

ΔHrxn = (ΔH bonds broken) + (ΔH bonds formed)

242
Q

What are some rules of entropy

A
  • S solid < S liquid < S gas

- all spontaneous processes increase the entropy of the system

243
Q

how do you find ΔS

A 
ΔS = qrev/T
ΔS = Sfinal-Sinitial
244
Q

When has entropy reached it maximum state

A

at equilibrium

245
Q

how do you find ΔS

A 
ΔS = qrev/T
ΔS = Sfinal-Sinitial
246
Q

what does Gibbs Free Energy tell you

A

the spontaneity of a reaction

247
Q

What is the equation for Gibbs Free energy

A

ΔG = ΔH - TΔS

248
Q 
What happens if ΔH and ΔS are...
\+   +
\+   -
-   +
-    -
A

+ + = spontaneous at high temperature
+ - = non spontaneous
- + = spontaneous at all temperatures
- - = spontaneous at low temperatures

249
Q

What does ΔG = when considering Keq

A 
ΔG = -RT ln Keq
ΔG = -RT ln Q
ΔG = ΔG + RT ln Q
250
Q

What does ΔG = when considering Keq

A 
ΔG = -RT ln Keq
ΔG = -RT ln Q
ΔG = ΔG + RT ln Q
251
Q

how many mmHG is 1 atm

A

760

252
Q

how many torr is 1 atm

A

760

253
Q

What are standard conditions

A

25 degrees celcius

198 K

254
Q

What is the equation for Boyle’s law (ideal gas)

A

P1V1 = P2V2

constant temperature

255
Q

What is the equation for charles law (ideal gas)

A

V1/T1 = V2/T2

constant pressure

256
Q

What is avogadros law (ideal gas)

A

n1/V1 = n2/V2

constant temp and pressure

257
Q

What is the ideal gas law

A

PV = nRT

258
Q

What are the two values for R in the ideal gas law

A

.0821 Latm/(molK)

8.314 J/(K*mol)

259
Q

How do you know which R to use in the ideal gas law

A

look at the units
.0821 (L*atm)
8.314 (J)

260
Q

How do you find density with the ideal gas law

A

n = m/MM
mass in grams over molar mass

so density = P(MM)/RT

261
Q

What is the volume of one mole of gas at standard temperature and pressure

A

22.4 L

262
Q

how are PV and T related

A 
P1V1/T1 = P2V2/T2
1= STP 
2 = actual conditions
263
Q

how do you find density of a gas with the ideal gas law

A 
1. find V2
V2 = 22.4 (P1/P2)(T2/T1)
2. d =m/V2
divide mass (g) by V2
264
Q

how do you find molar mass of a gas with the ideal gas law

A 
1. find Vstp
Vstp = V (P2/P1)(T1/T2)
2. find density
g/L (Vstp)
3. Molar mass = d*V1 (V1 = 22.4)
265
Q

What are causes of deviation to the ideal gas law

A
  1. High pressure

2. Low temperature

266
Q

What is daltons law of partial pressures

A

the total pressure of a gas is equal to the sum of the partial pressures in the gas

267
Q

how do you calculate the partial pressures of gasses with moles and pressure

A 
Pp = (P)(xa)
xa = number of moles of said gas/total moles of gas
268
Q

What are the assumptions of the kinetic molecular theory of gasses

A
  1. the size of particles in gasses are negligible compared to volume
  2. atoms and molecules don’t exhibit intermolecular attractions
  3. particles are in continuous random motion, and collisions
  4. collisions are elastic
  5. average KE is proportional to the average temperature
269
Q

What does KE for an ideal gas =

A

KE = 3/2 kT

270
Q

how are rates of diffusion related to molecular mass compared

A

r1/r2 = square root of MM2/MM1

271
Q

how are the rates of effusion related to molecular mass compared

A

the same way as diffusion

272
Q

what are amorphous solids

A

has no ordered arrangement of molecules

273
Q

What are the two kinds of crystals

A

metal and ionic

274
Q

what are the characteristics of ionic crystals

A
  • high boiling/melting points

- low conductivity

275
Q

What are the three types of unit cells

A

simple cubic
body-centered cubic
face centered cubic

276
Q

what happens to the liquid as some of it evaporates

A

the liquid cools due to a loss of kinetic energy

277
Q

what is the definition of boiling point

A

the temperature at which the vapor pressure of the liquid equals the external pressure

278
Q

what is the gibbs function for phase changes

A 
ΔG = G (gas) - G (solid) (if going from a gas to a solid)
os G(gas) = G(solid) at equilibrium
279
Q

Where is the gas phase usually found in a phase diagram

A

high temperatures, low pressures

280
Q

What is the liquid phase usually found in a phase diagram

A

high temperatures, high pressure

281
Q

Where is the solid phase usually found in a phase diagram

A

low temperatures, high pressure

282
Q

what is the triple point on a phase diagram

A

the point (pressure and temperature) where all three phases are in equilibrium

283
Q

What is the critical point

A

the point (temperature and pressure) at which no distinction between liquid and gas can be made

284
Q

what does adding a solute do to the freezing point of water

A

causes freezing point depression ( freezing point goes down)

285
Q

What is the equation for freezing point depression

A 
ΔT(f) = K(f)m
m = molality
K(freezing) = freezing point depression
286
Q

What is molality

A

moles/kg

287
Q

what happens to boiling point as you add solute to water

A

causes boiling point elevatoin

288
Q

What is the equation for boiling point elevation

A

ΔT(b) = K(b)m

289
Q

What is the equation for osmotic pressure

A 
II = MRT
II = osmotic pressure
M = molarity 
R = ideal gas constant
T = temperature (K)
290
Q

What is raoult’s law

A

P(a) = x(a)P(i)(a)
The vapor pressure of substance A is equal to the mole fraction of A multiplied by the vapor pressure of A (pure substance)

291
Q

What is solvation or dissolution

A

dissolving something in solution

292
Q

When does solvation occur

A

when attractive forces between the solute and solvent are greater than those between the solute particles

293
Q

what is the general rule of solvation

A

like dissolves like

294
Q

What is the solubility of a substance

A

the maximum amount of that substance that can be dissolved in a particular solvent and a particular temperature

295
Q

What happens when the maximum amount of a substance has been dissolved in a solvent

A

the solution is saturated

296
Q

What things are soluble in water

A

look at page 280 of kaplan

297
Q

What do the roman numerals after metal ions mead

A

they show how much of a positive charge the metal has

298
Q

What is ferr -ous and ferr -ic

A

ferrOUS means it is the lesser of two cations of iron

ferrIC means it is the greater of two cations of iron

299
Q

what is hydride, fluoride, oxide

A

-IDE means that they are monatomic anions

300
Q

what are oxyanions

A

anions that contain oxygen

301
Q

what is the difference between nitrite and nitrate

A

nitr -ITE is the lesser of two anions

nitr -ATE is the greater of two anions

302
Q

What do hypo and per mean when speaking of anions

A

Hypo- chlor -ite means it is the least of four anions
chrorite is the 3rd of four anions
chlorate is the 2nd of four anions
per- chlor -ate is the greatest of four anions

303
Q

what can be added in the name in place of hydrogen when one hydrogen has been added to an anion

A

bi-
like bicarbonate HCO3-
or bisulfate HSO4-

304
Q

when is a solute a strong electrolyte

A

if is completely dissolves into its constituent ions

305
Q

what is a weak electrolyte

A

a solute that doesn’t completely dissolve into its constituent ions

306
Q

what are non-electrolytes

A

solutes that don’t ionize at all in solution

307
Q

what is percent composition by mass

A

mass of the solute divided by mass of the solvent

multiplied by 100

308
Q

what is the mole fraction

A

the moles of a compound divided by the rest of the moles in the solution

309
Q

What is molarity

A

moles of solute per liter of solution (SOLUTION)

310
Q

What is molality

A

moles of solute per kilogram of solvent (SOLVENT)

311
Q

what is normality

A

the number of gram equivalent weights of solute per liter of solution (SOLUTION)

312
Q

what is the equation for dilution

A

M1V1 = M2V2
Molarity
Volume

313
Q

what is the ion product (I.P.) also referred to as the Qsp

A

IP = [An+]m[Bm-]n

314
Q

What is the solibility product constant

A

Ksp = [An+]m[Bm-]n (in a saturated solution/equilibrium solution)

315
Q

What is the difference in Qsp and Ksp

A

Qsp refers to the starting ion product

Ksp refers to the equilibrium ion product

316
Q

If Qsp > Ksp

A

precipitation will occur

317
Q

if Qsp = Ksp

A

nothing will change, the solution is at equilibrium

318
Q

if Qsp < Ksp

A

the solid will dissolve until saturation

319
Q

what does the common ion effect state

A

that a salt will not dissolve as well in a solution that already has a common ion

320
Q

what are indicators

A

things that change colors in acids and bases

321
Q

what is litmus paper

A

an acid/base indicator

322
Q

what is the arrhenius definition of an acid and a base

A 
acid = something that produces  an H
base = something that produces an OH
323
Q

What is the Bronstead lowry definition of an acid and a base

A 
acid = proton donor
base = proton acceptor
324
Q

what is the lewis definition of an acid and a base

A 
acid = electron acceptor
base = electron donor
325
Q

What is F- called

A

fluoride

326
Q

what is HF called

A

hydrofluoric acid

327
Q

what is Br- called

A

bromide

328
Q

what is HBr called

A

hydrobromic acid

329
Q

What is ClO- called

A

hypochlorite

330
Q

What is HClO called

A

hypochlorous acid

331
Q

What is ClO2- called

A

chlorite

332
Q

what is HClO2 called

A

chlorous acid

333
Q

what is ClO3- called

A

chlorate

334
Q

what is HClO3 called

A

chloric acid

335
Q

what is ClO4- called

A

perchlorate

336
Q

what is HClO4 called

A

perchloric acid

337
Q

what is No2- called

A

nitrite

338
Q

what is HNO2 called

A

nitrous acid

339
Q

what is NO3- called

A

nitrate

340
Q

what is HNO3 called

A

nitric acid

341
Q

what is pH

A

the measure of Hydrogen ion concentration

342
Q

what does pH =

A 
pH = -log [H+]
pH = log 1/[H+]
343
Q

what is pOH

A

the measure of hydroxide ion concentration

344
Q

what does pOH =

A 
pOH = -log[OH-]
pOh = log 1/[OH-]
345
Q

What is Kw

A

the water dissociation constant

346
Q

what does Kw =

A

Kw = [H+][OH-} = 10^-14

347
Q

pH + pOH =

A

pH + pOH = 14

348
Q

how do yo approximate Ka = 1.8 * 10^-5

A

Ka = 5 -log 1.8
if the -log number is less that 3.2 the number will be closer to 5 than 4 , if it is more that 3.2 the number will be closer to 4 than 5

349
Q

what does the log 10^x =

A

x

350
Q

what is Ka and what does it indicate

A

it is the acid dissociation constant, and it measures how much the acid dissociates

351
Q

What is the Kb and what does it indicate

A

it is the base dissociation constant, and it measures how much the base dissociates

352
Q

What does Ka =

A

Ka = [H3O+][A-]/[HA]

353
Q

what kind of Ka does a weak acid have

A

a small one, the weaker the acid, the smaller the Ka

354
Q

What does Kb =

A

Kb = [B+][OH-]/[BOH]

355
Q

if Ka is large then Kb is

A

small

356
Q

what is a neutralization reaction

A

when an acid and a base neutralize each other, forming a salt

357
Q

what is a hydrolysis reaction

A

when the salt ions react with water to form acid and base

358
Q

What happens when you add strong acid and strong base

A

you get a salt and water (pH = 7)

359
Q

what happens when you add strong acid and weak base

A

you get a salt, but no water. and the cation of the salt can react to reform the weak base, but decreases the ph by using up the OH-

360
Q

what happens when you add weak acid and strong base

A

you get a salt, the anion of the salt reacts to form the weak acid, the pH decreases because the strong base uses up the H+

361
Q

What happens when you add a weak acid and a weak base

A

if the Kb is stronger than the Ka than the solution is basic

362
Q

What does Ka =

A

Ka = [H3O+][A-]/[HA]

363
Q

what kind of Ka does a weak acid have

A

a small one, the weaker the acid, the smaller the Ka

364
Q

What does Kb =

A

Kb = [B+][OH-]/[BOH]

365
Q

if Ka is large then Kb is

A

small

366
Q

what is a neutralization reaction

A

when an acid and a base neutralize each other, forming a salt

367
Q

what is a hydrolysis reaction

A

when the salt ions react with water to form acid and base

368
Q

What happens when you add strong acid and strong base

A

you get a salt and water (pH = 7)

369
Q

what happens when you add strong acid and weak base

A

you get a salt, but no water. and the cation of the salt can react to reform the weak base, but decreases the ph by using up the OH-

370
Q

what happens when you add weak acid and strong base

A

you get a salt, the anion of the salt reacts to form the weak acid, the pH decreases because the strong base uses up the H+

371
Q

What happens when you add a weak acid and a weak base

A

if the Kb is stronger than the Ka than the solution is basic

372
Q

what is polyvalence

A

when an acid or base has more than 1 equivalent of H+ or OH-

373
Q

What is normality the quantity of acidic or basic capacity

A

the amount of mole equivalents an acid or base will give off

374
Q

What happens in the Henderson hasselbach equation when conjugate acid concentration = weak base concentration

A

pOH = pKb

we are halfway to the equivalence point

375
Q

what are amphoteric or amphiprotic species

A

species that can act as either an acid or a base

376
Q

what are the steps in balancing a redox reaction

A
  1. separate the two half reactions
  2. balance all atoms in both reactions (use H2O to balance O, then use H+ to balance the H atoms)
  3. add electrons to balance the charges
  4. multiply both sides by a number so they both have the same number of electrons
  5. add the half reactions, and cancel out items
  6. check to see that they are all balanced
377
Q

what is the equivalence point

A

when the acid equivalents = base equivalents

378
Q

what is the equivalence point for strong acid and strong base

A

pH = 7

379
Q

how do you calculate the volume added to reach endpoint in a titration

A

VN (acid) = VN (base)
V = volume
N = Normality

380
Q

what is a buffer

A

a mixture of weak acid and it’s salt or a weak base and it’s salt.

381
Q

what do buffers do

A

they resist changes in pH when acids or bases are added

382
Q

What is the henderson hasselbach equation for a weak acid buffer solution

A

pH = pKa + log (conjugate base/weak acid)

383
Q

What is the henderson hasselbach equation for a weak base buffer solution

A

pOH = pKb + log (conjugate acid/weak base)

384
Q

What happens in the Henderson hasselbach equation when conjugate acid concentration = weak base concentration

A

pOH = pKb

we are halfway to the equivalence point

385
Q

What are the rules for assigning oxidation numbers

A
  1. oxidation of free elements is zero
  2. oxidation of monatomic ion is = to the charge
  3. oxidation number of group IA = 1, IIA = 2
  4. oxidation number of grooup VIIA = -1 (except when paired with a more electronegative atom)
  5. oxidation of H = -1 in compounds with less electronegative elements than H
  6. Oxidation number of O = -2, unless paired to a more electronegative atom O = +1, or unless two are paired O = -1
  7. sum of oxidation numbers in a neutral compound = 0
386
Q

what are the steps in balancing a redox reaction

A
  1. separate the two half reactions
  2. balance all atoms in both reactions (use H2O to balance O, then use H+ to balance the H atoms)
  3. add electrons to balance the charges
  4. multiply both sides by a number so they both have the same number of electrons
  5. add the half reactions, and cancel out items
  6. check to see that they are all balanced
387
Q

what types of reactions occur in galvanic cells

A

spontaneous redox reactions

388
Q

What is the set up of a galvanic cell

A
  1. Two half cells
    - One with a negative anode, the other with a positive cathode
  2. the anode and cathode are connected to allow for electron flow
  3. The anode is in an aqueous common ion solution
  4. The cathode is also in a common ion solution
  5. there is a salt bridge with uncommon ions to both solutions
389
Q

how does a galvanic cell work.

A
  1. the - anode is oxidized, losing it’s positive ion, and sending electrons through the wire
  2. those electrons move across the wire to the positive cathode
  3. Those electrons on the cathode cause the positive ions in solution to bind to the cathode and the electrons in reduction.
  4. ions in the salt bridge come out to counteract the ions added to the anode solution, and replace the ions lost from the cathode solution
390
Q

What are the rules in drawing a cell diagram of a galvanic cell

A
  1. the reactants and products are always listed from left to right in this form
    anode / anode solution // cathode solution / cathode
  2. single verticle lines represent a phase boundry
  3. double lines represent a salt bridge
391
Q

What is the reaction like in an electrolytic cell

A

non-spontaneous redox reaction

392
Q

What is the reduction potential of a species

A

the tendency of a species to acquire electrons and be reduced

393
Q

The more positive the reduction potential of a species…

A

the more likely it is to be reduced

394
Q

what is the standard electromotive force (emf)

A

the difference in potential energy between two half cells

395
Q

how do you get emf

A

emf = reduction potential red. + reduction potential cat

396
Q

What is the sign for reduction potential

A

E (knot)

397
Q

what is the equation for determining spontaneity from emf

A

ΔG = -nFE(cell)
n = number of moles of electrons exchanged
F - faradays constant
E = emf of the cell

398
Q

What to remember about ΔG and emf

A

if faradays constant is in j/V, then ΔG must be in j not kj

399
Q

what is the nernst equation

A

Ecell = E(knot)cell - (RT/nF)(lnQ)

400
Q

what is another way to get ΔG(Knot)

A 
ΔG(knot) = -RT ln Keq
R = 8.314 J/ (k*mol)
401
Q

nFE(knot)cell = RT ln Keq

A

nFE (knot)cell = RT ln Keq

(if nFE is positive then Keq is positive and the reaction moves to the right

402
Q

what is the mass defect

A

the fact that every nucleus has a smaller mass than the combined masses of protons and neutrons

403
Q

what is the cause of mass defect

A

that some mass has been converted into energy (binding energy)

404
Q

what atom has the most binding energy

A

iron

405
Q

how do you find binding energy

A

find the mass defect and plug it into E = mc^2

406
Q

What is fusion

A

the combining of small nuclei to form a larger nucleus

407
Q

what is fission

A

the splitting of a large nucleus to from two smaller nuclei

408
Q

can cause fission

A

the addition of a neutron

409
Q

What is alpha decay

A

loss of a 4HE (2 protons, 2 neutrons)

410
Q

what happens in alpha decay

A
  • 4 mass number

- 2 atomic number

411
Q

What is beta decay

A

loss of an electron

412
Q

which is more penetrating alpha or beta decay

A

beta

413
Q

What is alpha decay

A

loss of a 4HE (2 protons, 2 neutrons)

414
Q

what happens in alpha decay

A
  • 4 mass number

- 2 atomic number

415
Q

What is beta decay

A

loss of an electron

416
Q

which is more penetrating alpha or beta decay

A

beta

417
Q

What are the two type of beta decay

A

B- (an electron)

B+ (a positron)

418
Q

What does a negative beta decay do

A

+1 in atomic number

419
Q

What does positron emission do

A

-1 in atomic number

420
Q

What is gamma decay

A

emission of high energy photons

421
Q

what happens in gamma decay

A

nothing to mass and atomic number, just a loss of energy

422
Q

what is electron capture

A

gaining an electron that combines with a neutron to form a proton

423
Q

what happens in electron capture

A

no change to mass number

+1 atomic number

424
Q

what is the equation for exponential decay

A 
n = n(knot) e^-(gamma t)
n = number that remain
n(knot) = starting number
gamma = (decay constant)
t = time
425
Q

What is avogadros number

A

6.02 x 10^23

DAT (14 decks)

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