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1
Q

7 Diatomic Molecules

A
  1. Hydrogen
  2. Nitrogen
  3. Oxygen
  4. Fluorine
  5. Chlorine
  6. Bromine
  7. Iodine
    - Should always have a 2 subscript when writing the formula for diatomic molecules because they need to bond even if its with itself
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2
Q

Rules for naming a covalent compound

A
  1. The first element is named first, using the the elements name
  2. The second elements is names as an anion (suffix-ide)
  3. Prefixes are used to denote the number of atoms
  4. “Mono” is not used to name the first element
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3
Q

Prefixes

A
  1. Mono
  2. Di
  3. Tri
  4. Tetra
  5. Penta
  6. Hexa
  7. Hepta
  8. Octa
  9. Nona
  10. Deca
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4
Q

Rules for naming Ionic Compounds

A
  1. Balance Charges (charges should equal zero)
  2. Cation is always written first (in name and in formula)
  3. Change the ending of the anion to -ide
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5
Q

Rules for naming an acid

A
  1. when the name of the anion ends in -ide, the acid name begins with the prefix hydro-, the stem of the anion has the suffix -ic, and it is followed by the word acid
  2. When the anion name ends in -ite, the acid name is the stem of the anion with the suffix -ous, followed by the word acid
  3. When the anion name ends in -ate, the acid name is the stem of the anion with the suffix -ic followed by the word acid
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6
Q

Oxidation number

A

A number assigned to an atom in a molecular compound or molecular ion that indicated the general distribution of electron among the bonded atoms

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7
Q

Determining oxidation number

A
  1. The oxidation number of any uncombined element is 0
  2. The oxidation number of a monoatomic ion equals the charge on the ion
  3. The more electronegative element in a binary compound is assigned the number equal to the charge it would have if it were an ion.
  4. The oxidation number of fluorine in a compound is always -1
  5. Oxygen has an oxidation number of -2 unless it is combined with F, when it is +2, or it is in peroxide, when it is -1
  6. The oxidation state of hydrogen in most of its compounds is +1 unless it combined with a metal, in which case it is -1
  7. In compounds, the elements of groups 1 and 2 as well as aluminum have oxidation number of +1, +2, and +3, respectively
  8. The sum of the oxidation numbers of all atoms in a neutral compound is 0
  9. The sum of the oxidation number of all atoms in a polyatomic ion equals the charge of the ion
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8
Q

Variable valences for transitional metals- metals

A
  • Chromium
  • Manganese
  • Iron
  • Cobalt
  • Copper
  • Lead
  • Mercury
  • Tin
  • Gold
  • Silver
  • Bismuth
  • Antimony
  • Cadmium
  • Zinc
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9
Q

Variable valences for Chromium

A
  • Cr
  • Charge: +2, +3
  • Stock name: Chromium (II), Chromium (III)
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10
Q

Variable valences for Manganese

A
  • Mn
  • Charge: +2, +3
  • Stock name: Manganese (II), Manganese (III)
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11
Q

Variable valences for Iron

A
  • Fe
  • Charge: +2, +3
  • Stock name: Iron (II), Iron (III)
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12
Q

Variable valences for Cobalt

A
  • Co
  • Charge: +2, +3
  • Stock name: Cobalt (II), Cobalt (III)
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13
Q

Variable valences for Copper

A
  • Cu
  • Charge: +1, +2
  • Stock name: Copper (I), Copper (II)
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14
Q

Variable valences for Lead

A
  • Pb
  • Charge: +2, +4
  • Stock name: Lead (II), Lead (IV)
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15
Q

Variable valences for Mercury

A
  • Hg
  • Charge: +1, +2
  • Stock name: Mercury (I) (Hg_2^+2), Mercury (II)
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16
Q

Variable valences for Tin

A
  • Sn
  • Charge: +2, +4
  • Stock name: Tin (II), Tin (IV)
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17
Q

Variable valences for Gold

A
  • Au
  • Charge: +1, +3
  • Stock name: Gold (I), Gold (III)
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18
Q

Variable valences for Silver

A
  • Ag
  • Charge: +1, +2
  • Stock name: Silver (doesn’t need a roman numeral because it is almost always 1+), Silver (II) (rarely)
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19
Q

Variable valences for Bismuth

A
  • Bi
  • Charge: +3, +5
  • Stock name: Bismuth (III), Bismuth (V)
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20
Q

Variable valences for Antimony

A
  • Sb
  • Charge: +3, +5
  • Stock name: Antimony (III), Antimony (V)
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21
Q

Variable valences for Cadmium

A
  • Cd
  • Charge: +2
  • Stock name: Cadmium (no roman numeral needed because charge doesn’t change)
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22
Q

Variable valences for Zinc

A
  • Zn
  • Charge: +2
  • Stock name: Zinc (no roman numeral needed because charge doesn’t change)
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23
Q

Ammonium

A

NH4, +1

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24
Q

Acetate

A

C2H302, -1

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25
Q

Bromate

A

BrO3, -1

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26
Q

Chlorate

A

ClO3, -1

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27
Q

Chlorite

A

ClO2, -1

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28
Q

Cyanide

A

CN, -1

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29
Q

Dihydrogen phosphate

A

H2PO4, -1

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30
Q

Hypochlorite

A

ClO, -1

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31
Q

Hydrogencarbonate (bicarbonate)

A

HCO3

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32
Q

hydrogen sulfate (bisulfate)

A

HSO4, -1

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33
Q

hydrogen sulfite (bisulfite)

A

HSO3, -1

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34
Q

hydroxide

A

OH, -1

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35
Q

iodate

A

IO3, -1

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36
Q

Nitrite

A

NO2, -1

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37
Q

Perchlorate

A

ClO4, -1

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38
Q

Permanganate

A

MnO4, -1

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39
Q

Thiocyanate

A

SCN, -1

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40
Q

Carbonate

A

CO3, -2

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41
Q

Chromate

A

CrO4, -2

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42
Q

Dichromate

A

Cr2O7, -2

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43
Q

Oxalate

A

C2O4, -2

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44
Q

Selenate

A

SeO4, -2

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45
Q

Silicate

A

SiO3, -2

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46
Q

Sulfate

A

SO4, -2

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47
Q

Sulfite

A

SO3, -2

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48
Q

Phosphate

A

PO4, -3

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49
Q

Phosphite

A

PO3, -3

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50
Q

NH4

A

ammonium, +1

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51
Q

C2H3O2

A

acetate, -1

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52
Q

BrO3

A

bromate, -1

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53
Q

ClO3

A

Chlorate, -1

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54
Q

ClO2

A

Chlorite, -1

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55
Q

CN

A

cyanide, -1

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56
Q

H2PO4

A

dihydrogen phosphate, -1

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57
Q

ClO, -1

A

hypochlorite

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58
Q

HCO3

A

hydrogencarbonate (bicarbonate), -1

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59
Q

HSO4

A

hydrogen sulfate (bisulfate), -1

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60
Q

HSO3

A

hydrogen sulfite (bisulfite), -1

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61
Q

OH

A

hydroxide, -1

62
Q

IO3

A

iodate, -1

63
Q

NO2

A

nitrite, -1

64
Q

ClO4

A

perchlorate, -1

65
Q

MnO4

A

permanganate, -1

66
Q

SCN

A

thiocyanate, -1

67
Q

CO3

A

carbonate, -2

68
Q

CrO4

A

chromate, -2

69
Q

Cr2O7

A

dichromate, -2

70
Q

C2O4

A

oxalate, -2

71
Q

SeO4

A

selenate, -2

72
Q

SiO3

A

silicate, -2

73
Q

SO4

A

sulfate, -2

74
Q

SO3

A

sulfite, -2

75
Q

PO4

A

phosphate, -3

76
Q

PO3

A

phosphite, -3

77
Q

Soluble compounds rules

A
  • dissolve
    1. All compounds of the alkali metals (group 1A) are soluble
    2. All salts containing NH4^+, NO3^-, ClO4-, and C2H3O2^- are soluble
    3. All chlorides, bromides, and iodides (salts containing Cl-, Br-, or I-) are soluble
    4. All sulfates (salts containing SO4^-2) are soluble
78
Q

Exceptions to soluble compound rules

A
  • Insoluble, precipitate
    1. Halides where Ag+, Pb2^+2, are the cations are insoluble (ex. AgCl)
    2. Sulfates with the cations Hg+2, Pb+2, Ca+2, Sr+2, and Ba+2 (ex. BaSO4)
79
Q

Insoluble Compounds rules

A
  • precipitate
    1. All hydroxides (OH- compounds) and all metal oxides
    (O-2 compounds) are insoluble.
    2. When metal oxides do dissolve they react with water to form hydroxides
    3. All compounds that conation PO4^-3, CO3^-2, SO3^-2, and S^-2 are insoluble
80
Q

Exceptions to insoluble rules

A
  • Soluble, dissolve
    1. Hydroxides and metal oxides combine with group 1A elements and Ca+2, Sr+2, and Ba+2
    2. Group 1A and NH4^+ compounds are soluble
81
Q

Describe how to put numbers into scientific notation and regular notation

A
  1. Into scientific notation
    - only one number before decimal
    - + if you made the number smaller, - if you made the number bigger
  2. Into regular notation
    - if +, move the decimal that many to the right
    - If -, move the decimal that many to the left
82
Q

Describe how to add and subtract numbers in scientific notation

A
  1. Determine the number by which to increase the smaller exponent by so it is equal to the larger exponent.
  2. Increase the smaller exponent by this number and move the decimal point of the number with the smaller exponent to the left the same number of places. (i.e. divide by the appropriate power of 10 .)
  3. Add or subtract the new coefficients.
  4. If the answer is not in scientific notation (i.e. if the coefficient is not between 1 and 10 ) convert it to scientific notation.
  • move the decimal of the number with the smaller exponent to the left
83
Q

Describe how to multiply numbers in scientific notation

A
  1. Multiply the coefficients–round to the number of significant figures in the coefficient with the smallest number of significant figures.
  2. Add the exponents.
  3. Convert the result to scientific notation.
84
Q

Describe how to divide numbers in scientific notation

A
  1. Divide the coefficients–round to the number of significant figures in the coefficient with the smallest number of significant figures.
  2. Subtract the exponents.
  3. Convert the result to scientific notation.
85
Q

counting significant digits

A
  • If the decimal point is absent, start counting from the atlantic (right) side, starting with the first non zero digit
  • If the decimal point is present, start counting from the pacific (left) side, starting with the first non zero digit
86
Q

Significant digit definition

A

Every number that is measured plus one that is estimates. Whenever you are using a measuring device, you should record your measurement so that the last number is estimates

  • we usually use scientific notation to avoid confusion about significant figures
87
Q

Define the rules for deciding the number of significant figures in a measurement

A

estimate one figure more than your measurement tool lets you

88
Q

Exact number

A

numbers that have defined values or are integers that result from counting numbers of objects, infinite significant digits
- Ex. counting digits and conversion factos (1000g in a kg)

89
Q

Adding and subtracting significant digits

A

find the guessed value in each number, after adding, round the answer at the first guessed place value

90
Q

multiplying and dividing significant digits

A

round to the same number of significant digits as the number with the least amount of significant digits

91
Q

rounding numbers

A
  • less than 5, round down

- 5 or greater round up

92
Q

Density

A
  • The mass of a substance per until volume of the substance
  • Density=mass/volume
  • Can be determined through the water displacement method. The object is massed then submerged in a measured amount of water in a graduated cylinder. The final volume in the graduated cylinder is read. The volume displaced by the object is the volume of the object.
93
Q

accuracy

A

the agreement of a particular value with a true value

94
Q

precision

A

the degree of agreement among several measurements of the same quantity. The degree of precision refers to the number of digits that a measuring device permits one to measure. In a measuring divide, all except the last digit, which is estimated, are certain.

  • For example, a balance which measures to the nearest .0001 g is more precise than one that measures to the nearest .001 g.
95
Q

random error

A

(indeterminate error) means that a measurement has an equal probability or being high or low

96
Q

systematic error

A

(determinate error) occurs in the same direction each time; it is either always high or always low

  • ex. A balance could have a defect that caused it to read 1.000 g too high every time. A thermometer could have a defect that caused it to read 1.00 degrees C too low every time
97
Q

Percent error=

A

Experimental value - Actual value X100 over actual value

98
Q

Temperature in science

A
  • science uses Kelvin
  • no negative numbers
  • 0= matches the motion of no particle motion, particle motion stops
  • When kinetic energy is zero, the temperate is as low as it can possible go
99
Q

temperature

A

A measure of the average kinetic energy (energy of motion) of a substance

100
Q

To convert celsius to Kelvin

A

K= C + 273

(°F - 32) x 5/9 = °C

101
Q

Solid

A
  • Shape: Definite
  • Volume: Indefinite
  • Particle Motion: Slow- vibrate in place
  • Particle attractions: strongest attraction of the 3 states OR particles are moving slow enough that they cannot overcome attractions
102
Q

Liquid

A
  • Shape: Indefinite
  • Volume: Definite
  • Particle Motion: Particles can move past each other but are moving slow enough to experience some attraction
  • Particle attractions: Particles are attracted to each other but can overcome the attraction to move past each other. They cannot overcome the attractions enough to escape the liquid entirely.
103
Q

Gas

A
  • Shape: Indefinite
  • Volume: Indefinite
  • Particle Motion: particles moving very fast. They move so fast that particle cannot be impacted by attractions.
  • Particle attractions: Particle attractions are irrelevant until there is a high pressure or low temperature
104
Q

Substance vs. Mixture

A
  • substances can be elements or compounds,

- mixtures are the combination of elements and/or compounds

105
Q

Heterogeneous mixture

A
  • Particles are not small enough to the point where they are invisible/solution is clear
106
Q

Homogeneous mixture

A
  • Solution
  • Very small particles that are equally/evenly distributed
  • Ex. Coolaid
  • 2 parts to a solution- Solvent and solute
  • Solvent- bigger part that does the dissolving
  • Solute- smaller part that gets dissolved
  • Soda is a solution when flat
  • Doesn’t have to be a solid and liquid- can be any phase. Ex. Alloy- 2 metals- pewter, brass, and bronze
107
Q

Physical changes

A

changes that do not change the original composition of the substance. Changes in state such as boiling or melting are physical changes. Changes involving an alteration in the form of the substance such as grinding or tearing are physical changes. Bonds are not broken and no reaction occurs.

  • Physical properties- properties that can be observed without changing the composition of the substance. Ex. Density, color, and boiling point
108
Q

Chemical change

A

Changes the composition of the original substance by breaking and making bonds between atoms. A new substance is produced when a chemical change occurs.

  • Evidence a chemical change has occurred includes: Change in color, change in odor, production of gas or solid (precipitate), change in energy. These indications do not always mean a chemical change has occurred, but they often do. Examples of chemical properties include: Flammability and reactivity to air
109
Q

Atom

A

smallest piece of an element

110
Q

molecule

A

smallest piece of a molecular compound (2 nonmetals)

111
Q

Types of Heterogeneous mixtures

A
  1. Colloid- Some medium sizes particles- light bounces or is bent when it hits the particles- the creates color, Ex. Mayonnaise- oil, vinegar, and egg clear, mixture becomes white.
  2. Suspension- Even larger particles that can be seen with the naked eye, Many times, particles are so large, they settle with gravity, Ex. Italian dressing, OJ with pulp, sand in water, Easy to separate
112
Q

Types of Homogeneous mixture

A

Solution- Very small particles that are equally/evenly distributed

113
Q

Matter chart

A

see chart

114
Q

Finding the electron

A
  • JJ Thompson used the cathode ray tube and proved the presence of electrons and developed the plum pudding model/sea of electrons.
  • Robert Millikan used the oil drop experiment to determine the mass and charge of an electron
115
Q

Finding the proton

A

Ernest rutherford discovered protons through his gold foil experiment when some of the positively charged gamma particles were reflected, proving that there were positive centers in an atom.

116
Q

Finding the nucleus

A

Ernest rutherford used the gold foil experiment to determine that atoms had a nucleus, and the nucleus was made out of positive particles called protons, and that the atom is mostly empty space.

117
Q

Periodic table developments

A
  1. Johann Dobereiner- German chemist, foreshadowed the periodic for the chemical elements, discovered trends in certain properties of select groups of elements, proposed law of triads.
  2. John Newlands- Devised the first periodic table arranged in order of their relative atomic weights, law of octaves
  3. Lothar Meyer- periodic classification of the elements, organized 28 elements by atomic weight and valence electrons, possible inspired Mendeleev
  4. Dmitri Mendeleev- came to him in a dream, proposed an organized scheme of all the known elements based on valence electrons and atomic weight, predicted the existence of 8 new elements and their masses
  5. Henry Moseley- determined that each element had a unique atomic number using x-rays
118
Q

Protons

A
  • subatomic particle with a positive charge, found in the nucleus of an atom, amount indicated by atomic number
  • subatomic particle that identifies the element
119
Q

neutrons

A
  • subatomic particle with a neutral/no charge, found in the nucleus of the atom
120
Q

electrons

A
  • subatomic particle with a negative charge, found around the nucleus of the atom
121
Q

Ion

A

an atom, radical, or molecule that has lost or gained one or more electrons and has a positive (cation) or negative (anion) charge

122
Q

isotopes

A

An atom that has the same number of protons (or the same atomic number) as other atoms of the same element do, but that has a different number of neutrons (and thus a different atomic mass).
- Ex. C-14 and C-12

123
Q

atomic number indicates

A

number of protons

124
Q

Negative charge means

A

additional electrons were gained by the atom

125
Q

Positive charge means

A

electrons were lost by the atom

126
Q

determining number of P, N, E

A
  • Neutrons=mass number - atomic number
  • Protons= atomic number
  • Electrons= protons if the charge of the atom is 0
127
Q

Reading element on periodic table

A

atomic number over atomic mass

128
Q

When writing an elemental symbol

A

mass number/atomic number, symbol - opposite of periodic table

129
Q

How to know if you have an ionic compound

A

Contains metal and nonmetal or polyatomic ions

130
Q

How to know if you have a covalent compound

A

2 nonmetals or metalloid and nonmetal

131
Q

How to know if you have and acid

A

starts with H (except for water)

132
Q

molecular is the same as

A

covalent

133
Q

balanced chemical equation

A
  • uses coefficients to indicate the same number of atoms on both sides of the equation
  • true because of the law of conservation of mass
134
Q

mole

A

a relative unit used to mass a substance. I mole of anything=6.02X10^23

135
Q

avogadros number

A
  • 6.02 x 10^23 particles per mole.
  • the number of particles found in one mole of a substance
  • It is the number of atoms in exactly 12 grams of carbon-12.
136
Q

molar mass

A
  • grams per mole of a substance.
  • find by multiplying the atomic mass of each element in the compound by the number of particles in that compound and adding them together
137
Q

molar volume

A
  • The volume occupied by one mole of a material.

- the molar volume of an ideal gas at STP is 22.4 L/mol.

138
Q

particles can include

A

ions, atoms, molecule, or formula units

139
Q

FU vs. molecule

A
  • A FU is ionic, if broken down it consists on ions

- a molecule is covalent, if broken down it consists of atoms

140
Q

Old general chemistry conversion factors

A
  1. Given value / 1
  2. molar mass / 1 mol __
  3. avogadros # particle / 1 mol __
141
Q

when counting atoms and ions

A

subscripts of whole thing. so for polyatomic ions, subscript of the whole polyatomic ion, not each individual element in it

142
Q

limiting reagent

A

The reactant in a chemical reaction that limits the amount of product that can be formed. The reaction will stop when all of the limiting reactant is consumed.

143
Q

excess reagent

A

The reactant in a chemical reaction that remains when a reaction stops when the limiting reactant is completely consumed. The excess reactant remains because there is nothing with which it can react.

144
Q

steps for finding limiting reagent

A
  1. convert both reactants to moles of the same product
  2. compare amount of product produced from each
    - – the lesser quantity of products was produced by the limiting reactant
  3. convert from the given amount of the limiting reagent to the desired answer
145
Q

empirical formula

A

the lowest whole number ratio of elements in a compound

146
Q

how to find the empirical formula

A
  1. If you are given a percent, then assume 100%=100 grams. If you are given grams, start at #2. If you are given moles, start at #3
  2. convert from grams to moles using molar mass
  3. Divide all numbers of moles by the smallest number of moles to get the ratio of moles of each element
  4. If the numbers are whole numbers, use them as subscripts when you write your formula
  5. If the numbers are not whole numbers, multiply all numbers by an appropriate integer to obtain whole numbers. Then use the new, whole numbers as subscripts when you write the formula
147
Q

Molecular formula

A

includes every element actually present in the compound

148
Q

how to find molecular formula

A
  1. Find empirical formula if it was not already given
  2. Find the molar mass of the empirical formula
  3. Divide the molar mass of the molecular formula by the molar mass of the empirical formula
  4. Multiply the subscripts of the empirical formula by the number obtained in number 3.
  5. Write your new formula
149
Q

Predicting prodcuts

A

retype

150
Q

New conversion factors

A
  1. 22.4L/1 mol gas
  2. # total atoms/ions (subscripts) of separate elements in a compound / 1 molecule/FU of compound
  3. # moles of 1 element in a compound (subscript) / 1 mole compound
  4. # of moles of one element/compound in the equation (coefficients)/ # of moles of a noter element/compound in the equation (coefficients)

AP Chemistry (23 decks)

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