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Flashcards in Chemistry Deck (264)
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1
Q

naming the acid when it has only an element following the H

A

use the prefix hydro-, followed by the element’s root name and an-ic ending

2
Q

If the acid has an-ate polyatomic ion after the H

A

that makes it an-ic acid. H2SO4 is sulfuric acid

3
Q

When the acid has an-ite polyatomic ion after the H

A

makes it an-ous acid

4
Q

Atoms make bonds because

A

of their closed shell electron configurations

5
Q

why can’t nobel gases accept another electron into the shell

A

Because the valence shell of a noble gas is completely full

6
Q

the loss of an electron from a noble gas is unfavorable because

A

The nucleus is positively charged and pulls on the electrons

7
Q

Noble gases are

A

inert/unreactive

8
Q

Any element other than a noble gas has

A

an open shell configuration, which is unstable relative to the configuration of a noble gas.

9
Q

Non-noble atoms react to form bonds in an attempt to

A

achieve a closed shell electron configuration

10
Q

Even though the reaction may appear to be favorable because of its production of a closed shell species

A

there is a way to have both F atoms achieve a noble gas configuration.

11
Q

covalent bonds also occur between

A

diatomics

12
Q

sharing of electrons is called

A
  • covalent bond

- 2 nonmetals

13
Q

When a highly electronegative atom and an electropositive one are bonded together

A

an electron is transferred from the electropositive atom to the electronegative atom to form a cation and an anion, respectively.

14
Q

how does distance affect attraction between the two ions

A

At large distances, there is a negligible energy of attraction between the two ions, but as they are brought closer together, they are attracted to one another.

15
Q

how does distance affect ion attraction?

A

Ions are actually repelled at small distances. To explain this observation, remember that the ions’ nuclei are both positively charged. When the nuclei approach each other, they repel strongly–accounting for the steep rise in potential as the ions get closer than the bond length.

16
Q

Ionic compounds are a part of

A

a crystal lattice–a three dimensional regular array of cations and anions.

17
Q

why do Ionic compounds form lattices?

A

due to the contributing coulombic attractions of having each cation surrounded by several anions and each anion surrounded by several anions.

18
Q

why are covalent bonds stable?

A
  • due to the build-up of electron density between the nuclei.
  • By sharing electron pairs nuclei can achieve octets of electrons in their valence shells, which leads to greater stability.
19
Q

why does a Lewis structure only counts valence electrons

A

because these are the only ones involved in bonding. (To calculate the number of valence electrons, write out the electron configuration of the atom and count up the number of electrons in the highest principle quantum number.)

20
Q

We can create bonds by having

A

two atoms come together to share an electron pair.

21
Q

how is a bonding pair of electrons is distinguished from a non-bonding pair

A

by using a line between the two atoms to represent a bond (A lone pair is what we call two non-bonding electrons localized on a particular atom.)

22
Q

In atoms, electrons reside in

A

orbitals of differing energy levels such as 1s, 2s, 3d, etc.

23
Q

orbitals represent

A

the probability distribution for finding an electron anywhere around the atom.

24
Q

Molecular orbital theory

A

posits the notion that electrons in molecules likewise exist in different orbitals that give the probability of finding the electron at particular points around the molecule.

25
Q

bond order

A
  • the number of bonds between atoms in a molecule.
  • The bond order is the difference in the number of electron pairs occupying an antibonding and a bonding molecular orbital.
26
Q

Steps for Basic Stoichiometry Calculations

A

4 Steps:

  • Balance the equation
  • Convert units of given substance to moles
  • Find moles of wanted substance using mole ratio
  • Convert moles of wanted substance to desired units
27
Q

Moles =

A

grams/formula mass

28
Q

Converting between gas volume and moles

A

Number of moles = PV/RT

29
Q

At STP, a mole of gas will always occupy

A

22.4 L of volume

30
Q

Avogadro’s Number

A
  • provides the conversion factor for moving from number of particles to moles.
  • There are 6.02×1023 formula units of particles in every mole of substance
31
Q

Limiting Reactant Problems

A
  • Whichever reactant that limits the production of product is the limiting reactant
  • Start by converting to molesO2(g)
  • Figure out how many grams of C(s) react. Mole ratio tells us that an equal number of moles will react.
  • Subtract to figure out how much C(s) is left over.
  • limiting reagent limits or determines the amount of product that can be formed.
32
Q

Percent Yield

A
  • (actual yield/theoretical yield)*100
  • ## How many grams of CaO should be produced? First verify the equation is balanced. Now convert to moles, based on the amount ofCaCO3 present, then to grams
33
Q

Energy Changes

A
  • ## Reactions that release energy in the form of heat are called exothermic reactions. Conversely, reactions requiring heat energy are known as endothermic reactions.
34
Q

Based on the following balanced equation, how many kcal of energy are needed to decompose 300 grams of CaCO3(s) ?

CaCO3(s) + 176 kJ→CaO(s) + CO2(g)

A
  • Convert to moles.
  • 300g/100g ×1 mole CaCO3 = 3 moles CaCO3
  • Now do the mole ratio; 176 kJ are needed for every mole of CaCO3 .
  • 3 molCaCO3/1 mol CaCO3 ×176 kJ = 528 kJ
  • put the answer in terms of kcal. There are 4.18 kJ in every kcal. Use the necessary conversion factor.
  • ## 528kj x 1.00kcal / 4.18kj = 126 kcal
35
Q

Enthalpy

A

how much heat a substance has at a given temperature and pressure, and is symbolized by the symbol H

36
Q

heat of reaction

A
  • The change in enthalpy for a reaction

- has symbol δH

37
Q

enthalpy determines

A
  • exothermic or endothermic
  • δH is negative for all exothermic reactions
  • positive for all endothermic reactions
38
Q

Heat of formation equation

A

δH = δH f (of all products) - δH f (of all reactants)

39
Q

Pressure=

A

Force/Area

40
Q

Boyle’s Law

A
  • Boyle’s law: P 1 V 1 = P 2 V 2
  • downward curve on volume v pressure graph
  • The most important thing to remember about Boyle’s Law is that it only holds when the temperature and amount of gas are constant.
  • A state of constant temperature is often referred to as isothermal conditions.
  • temp doesn’t have to be in kelvin
41
Q

Isothermal conditions

A

A state of constant temperature

42
Q

The Manometer

A
  • There are two ends to Boyle’s manometer. One end is open to the atmosphere. The other end is sealed, but contains gas at atmospheric pressure. Since the pressure on both ends of the tube is the same, the level of mercury is also the same.
43
Q

Manometer J-shaped, closed on curved end with V-50ml and open on straight end (Patm)

A
  • The pressure of the gas before mercury is added is equal to the atmospheric pressure, 760 mm Hg
  • After added mercury, the volume of the gas, V 2 , drops to 50 mL
  • P 2=P 1 V 1/V 2=
    (100 mL)(760 mm Hg)/(50 mL)=1520 mm Hg
44
Q

Charle’s Law

A
  • V1/T1 = V2/T2
  • If temperature is measured on a Celsius scale, T can be negative. The standard absolute scale is the Kelvin (K) scale. The temperature in Kelvin can be calculated via T k = T C + 273.15 .
  • straight line up on volume v temp graph
45
Q

Density=

A

PM/RT (where M is molar mass)

46
Q

Partial Pressure=

A
  • P tot = P A + P B + P C + …

- moles g A/total mols = PP gas A/total pressure

47
Q

Each individual gas obeys the ideal gas law so

A
  • we can rearrange PV= nRT to find pressure:

- P1=n1RT/V

48
Q

Arrhenius

A

defined acids to be proton (H+) donors and bases to be hydroxide ion (OH-) donors in aqueous solution.

49
Q

Bronsted-Lowry

A
  • describing acids as proton donors and bases as proton acceptors.
  • The Bronsted-Lowry model implies that there is a relationship between acids and bases (acids transfer protons to bases) and allows defines conjugate acids and conjugate bases
  • HA (acid) + B (base)= A-(CB) + BH+(CA)
50
Q

Lewis model

A

proposes that an acid is an electron pair acceptor while a base is an electron pair donor.

51
Q

pH =

A
  • log [H+]
52
Q

pOH

A

the negative common logarithm of the concentration of OH-

53
Q

pH + pOH =

A

14

54
Q

pH, pOH, [H+], and [OH-] at extreme acid

A
  • pH=14
  • pOH=0
  • [H+]=10^-14
  • [OH-]=10^0=1
55
Q

When a strong acid or a strong base is added to water

A

it nearly completely dissociates into its ion constituents because it has a pKa or pK b less than zero. For example, a solution of H2SO4 in water contains mostly H+ and SO4 2

56
Q

The concentration of acid equals

A

the concentration of H+

57
Q

Common strong acids that should be memorized

A
  • Hydrochloric HCl
  • Hydrobromic HBr
  • Hydroiodic HI
  • Nitric HNO3
  • Sulfuric H2SO4 (only strong first time)
  • Perchloric (HClO4)
58
Q

Strong bases include

A

Group I hydroxides (LiOH, NaOH, KOH, etc.) and Group II hydroxides except for Be(OH)2 and Ba(OH)2.

59
Q

why is calculating the pH of weak acid and weak base solutions is much more complicated than strong?

A

weak acids and bases do not completely dissociate in aqueous solution but are in equilibrium with their dissociated forms.

60
Q

calculate the pH of a 0.10 M solution of acetic acid in water

A

DEVANIKKKKK

61
Q

Polyprotic Acids

A

acids that can donate more than one proton per molecule.

62
Q

Two key features of polyprotic acids

A
  • they lose their protons in a stepwise manner

- each proton is characterized by a different pK a.

63
Q

factors contributing to the pK a of each acidic proton in a polyprotic species

A

The factors contributing to the pK a of each acidic proton in a polyprotic species are the same factors that determine the relative acidity of monoprotic acids–the dominant factor is strength of the acid-H bond.

64
Q

Buffers

A
  • A buffer is simply a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid
  • Buffers work by reacting with any added acid or base to control the pH.
  • A buffer works by replacing a strong acid or base with a weak one
65
Q

Calculating the pH of Buffered Solutions

A

Use the Henderson-Hasselbalch Equation

66
Q

Henderson-Hasselbalch Equation

A
  • pH=pKa+log(base/acid)
  • pH= PKa + log (A-/HA+)
  • POH=PKb + log (HA/A)
  • Find pKa using the expression that’s similar to equilibrium
67
Q

Titrations

A
  • A titration curve is drawn by plotting data attained during a titration, titrant volume on the x-axis and pH on the y-axis.
  • equivalence point at steepest part
68
Q

Redox Reaction

A

The oxidation state of an atom in a covalent compound is an imaginary charge assigned to that atom if all the electrons in its bonds were completely given to the more electronegative atom in the bond.

69
Q

redox rules

A
  1. Atoms in elemental form have oxidation states of zero.
  2. The charge on a monoatomic ion is equivalent to its charge.
  3. Hydrogen has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals.
  4. F, because it only forms one bond and is the most electronegative element, has an oxidation state of -1.
  5. O, unless bonded to F or itself, has an oxidation state of -2.
  6. The sum of all oxidation states in a compound must equal the total charge on the species. For hydrogen peroxide, we have seen that both O’s have -1 oxidation states and both H’s have +1 oxidation states, the sum of which is zero–the charge on hydrogen peroxide.
70
Q

Balancing Redox Reactions

A
  1. Separate oxidation and reduction half-reactions:
  2. Balance all atoms except for hydrogen and oxygen in each half-reaction. In this example they are already balanced.
  3. Balance oxygen by adding H2O as needed:
  4. To balance hydrogen, add H+ as needed: (Note: You still do this if you are in basic solution, later on, you will add OH- to “neutralize” the acid.)
  5. Balance the charge of each reaction by adding electrons to side with the greater charge:
  6. Multiply each half-reaction by the least integer factor that equalizes the number of electrons in each half-reaction. Then, add the half-reactions to obtain the overall balanced reaction in acidic solution:
  7. If your redox reaction is in acidic solution, the above reaction is properly balanced. However, if the reaction you wish to balance is in basic solution, you need to add these three steps:
  8. If the redox reaction is one in basic solution, then add OH- to both sides of the equation to “neutralize” each H+:
  9. “React” H+ and OH- to form H2O and eliminate water molecules on both sides of the equation
  10. Make sure that all atoms and charges are indeed balanced in your overall balanced equation for the redox reaction in basic solution
71
Q

Galvanic Cells

A
  • Galvanic cells harness the electrical energy available from the electron transfer in a redox reaction to perform useful electrical work. The key to gathering the electron flow is to separate the oxidation and reduction half-reactions, connecting them by a wire, so that the electrons must flow through that wire.
72
Q

two typical setups for galvanic cells

A
  • The left hand cell diagram shows and oxidation and a reduction half-reaction joined by both a wire and a porous disk
  • The right hand cell diagram shows the same cell substituting a salt bridge for the porous disk.
73
Q

Why is the salt bridge/porous disk necessary?

A

The salt bridge or porous disk is necessary to maintain the charge neutrality of each half-cell by allowing the flow of ions with minimal mixing of the half-cell solutions. As electrons are transferred from the oxidation half-cell to the reduction half-cell, a negative charge builds in the reduction half-cell and a positive charge in the oxidation half-cell. That charge buildup would serve to oppose the current from anode to cathode– effectively stopping the electron flow

74
Q

reduction/oxidation in relation to anode/cathode

A
  • reduction takes place at the cathode and oxidation takes place at the anode.
  • “The Red Cat ate An Ox”
75
Q

anode/cathode marked

A

The anode, as the source of the negatively charged electrons is usually marked with a minus sign (-) and the cathode is marked with a plus sign (+).

76
Q

Negatively charged electrons flow in a wire. Therefore

A

chemists indicate the direction of electron flow on cell diagrams and not the direction of current.

77
Q

cell potential, Ecell of a galvanic cell is measured in

A

volts

78
Q

how do we measure the Eo of any half- reaction.

A

by arbitrarily assigning a value of exactly zero to the potential of the standard hydrogen electrode

79
Q

E Cell example

A

DEVANIKK

80
Q

Electrolytic cells

A

like galvanic cells, are composed of two half-cells–one is a reduction half-cell, the other is an oxidation half-cell.

81
Q

the direction of electron flow in electrolytic cells

A

the direction of electron flow in electrolytic cells may be reversed from the direction of spontaneous electron flow in galvanic cells, the definition of both cathode and anode remain the same–reduction takes place at the cathode and oxidation occurs at the anode.

82
Q

comparing a galvanic cell to its electrolytic counterpart

A

When comparing a galvanic cell to its electrolytic counterpart, as is done in , occurs on the right-hand half-cell. Because the directions of both half-reactions have been reversed, the sign, but not the magnitude, of the cell potential has been reversed.

83
Q

The electrolytic cell reaction is not the only one occurring in the system-

A
  • the battery is a spontaneous redox reaction.
  • By Hess’s Law, we can sum the ΔG of the battery and the electrolytic cell to arrive at the ΔG for the overall process. As long as that ΔG for the overall reaction is negative, the system of the battery and the electrolytic cell will continue to function. The condition for ΔG being negative for the system (you should prove this for yourself) is that Ebattery is greater than - Ecell.
84
Q

Hess’s Law

A

regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes. This law is a manifestation that enthalpy is a state function.

85
Q

Electrolysis of Water

A
  • That spontaneous direction of reaction can be used to create water and electricity in a galvanic cell (as it does on the space shuttle). However, by using an electrolytic cell composed of water, two electrodes and an external source emf one can reverse the direction of the process and create hydrogen and oxygen from water and electricity.
  • The reaction at the anode is the oxidation of water to O2 and acid while the cathode reduces water into H2 and hydroxide ion.
  • requires input of energy because the products possess more chemical PE than H2)
86
Q

The Electrolysis of Water is

A

The electrolysis of pure water is -1.23 V. To make the electrolysis of water occur, one must apply an external potential (usually from a battery of some sort) of greater than or equal to 1.23 V. .

87
Q

Gibbs free energy

A
  • Gibbs free energy
  • DeltaH-TdeltaS
  • -G=spontaneous forward reaction
  • +G= spontaneous reverse reaction
  • G=0=equalibrium
  • DETERMINES SPONTANEITY
  • According to Hess’s Law, only state functions (G, H, S) and not path functions (w, q, E) of a series of reactions may be summed to generate a new value for the overall reaction.
  • You can add the ΔG or the ΔH of reactions using Hess’s Law
88
Q

When the temperature of a 20g sample of water is increased from 10C to 30C, the heat transferred to the water is

A
  • q=mcdeltaT
  • c is about 1cal/g
  • so 20g=20C
  • (20g)(1cal)(20C)=400
89
Q

An aqueous solution with pH5 at 25C has a hydroxide ion (OH-) concentration of

A

1x10^-9 molar

90
Q

The volume of water vapor required to produce 44.8 liters of oxygen by the below reaction is

2H2O(g)–>2H2(g)+O2(g)

A
  • 2 moles H2O to 1 mole O2
  • volume of gas will also be 2:1
  • so 89.6 liters H20 are required to produce 44.8 liters O2
91
Q

When 190 grams of MgCl2 are dissolved in water and the resulting solution is 500 ml in volume, what is the molar concentration of MgCl2 in the solution?

A
  • M=mol/l
  • find mass of 1 mol of MgCl2
  • determine moles of MgCl2 by dividing the mass of the sample by the molar mass of the compound (190/95.4=2 mol
  • the mol/v= 2/.5=4M
92
Q

a 600 ml container holds 2 moles of O2(g), 3 moles of H2(g), and 1 mole of He(g). Total pressure within the container is 760 torr. What is the partial pressure of O2?

A
  • mol gas a/total mols=pp gas a/total pressure
  • 2/6=x/760
  • x=253
93
Q

An ideal gas has a volume of 10 liters at 20C and a pressure of 750mmHg. Whats the volume of the same gas at STP?

A

PV=nRT

  • initial PV/T=final PV/T
  • (750)(10)/293=760V/273
  • 10 X 750/760 X 273/293
94
Q

When 3 moles of Fe2O3 are allowed to completely react with 56 grams of CO according to the above equation, approximately how many moles of iron, Fe, are produced?

Fe2O3(s)+3CO(g)–>2Fe(s)+3CO2(g)

A
  • find the limiting reactant
  • 56g CO=2 mol CO
  • Fe2O3:CO=1:3
  • CO is limiting
  • CO to Fe is 3:2, so 2 moles produces aout 1.3 moles of Fe
95
Q

organic compounds properties

A
  • organic compounds are much more soluble in nonpolar solvents
  • organic compounds don’t dissociate in solution
  • poor conductors of electricity
96
Q

isomers

A

same chemical makeup with identical constituent elements, but are arranged in a different geometrical arrangement,a and have different chemical properties

97
Q

hydrocarbons

A
  • the simplest organic compounds
  • contain only carbon and hydrogen
  • 3 categories:
    1. alkanes- contain only single C-C bonds
    2. Alkenes- contain double C-C bonds
    3. Alkynes- contain triple C-C bonds
98
Q

Alkanes

A
  • saturated hydrocarbons because each carbon atom is bonded to as many other atoms as possible
  • formula based on prefixes (CnH2n+2)
99
Q

prefixes

A
  • indicate the number of carbons in the hydrocarbon
  • meth=1
  • eth=2
  • prop=3
  • but=4
  • pent=5
  • hex=6
100
Q

Alkenes

A
  • unsaturated hydrocarbons because each carbon atom is not bonded to as many atoms as possible
  • Formula (CnH2n)
101
Q

suffixes

A
  • indicate what types of C-C bonds are present
  • ane=single bonds
  • ene=at least 1 double bond
  • yne=at least 1 triple bond
102
Q

alkynes

A
  • unsaturated hydrocarbons

- formula (CnH2n-n)

103
Q

hydrocarbon rings

A

aromatic hydrocarbons, the simplest of which is benzene C6H6

104
Q

Alcohol

A
  • an O-H bonded to a carbon atom. Because of hydroxyl (OH-) group, alcohols are polar
  • ol
  • ex. Methanol
105
Q

Halides

A
  • Halogen bonded to a carbon atom
  • can by named by suffixes: fluoride, chloride, bromide, iodide, or by prefixes: fluoro, chloro, broom,
  • ex. chloromethane
106
Q

organic acids or carboxylic acids

A
  • a carboxyl group (COOH) bonded to a carbon chain; the H dissociates, making these molecules acidic
  • ic acid
  • ex. acetic acid
107
Q

amines

A
  • an amine (NH2) bonded to a carbon atom; the amine group is similar to ammonia (NH3), and thus these molecules are basic
  • amine
  • ex. methylamine
108
Q

aldehydes

A
  • a carbonyl (C=O) group bonded to a terminal carbon, or a carbon that is at the end the carbon chain
  • aldehyde
  • ex. acetaldehyde
109
Q

ketones

A
  • a carbonyl (C=O) group that is bonded to a nonterminal carbon, or one that is not at the end of the carbon chain
  • one
  • ex. acetone
110
Q

Ethers

A
  • an oxygen atom links two hydrocarbon chains
  • ether
  • ex. dimethyl ether
111
Q

esters

A
  • an ester is essentially a ketone and an ether put together
  • ate
  • ex. methyl formate
112
Q

organic reactions

A
  • addition
  • substitution
  • polymerization
  • cracking
  • oxidation
  • esterification
113
Q

addition

A

a c-c double bond is converted into a single bond, freeing each of the two carbons to bond with another element. Also in addition reaction, a triple bond can be converted into a double bond

114
Q

substitution

A

one atom or group in a compound is replaced with another atom or group. chemically ,this very rarely happens

115
Q

polymerization

A

two smaller compounds, called monomers, are joined to form a larger third compound. in condensation polymerization, two monomers are joined in a reaction that produces water

116
Q

cracking

A

a larger compound is broken down into smaller compounds

117
Q

oxidation

A

an organic compound can react with oxygen at high temperatures to form carbon dioxide and water. This reaction should be familiar to you as combustion or burning

118
Q

esterification

A

an organic acid reacts with an alcohol to produce an ester and water

119
Q

four major groups of biomolecules

A
  • lipids
  • carbohydrates
  • nucleic acids
  • proteins
120
Q

lipids

A
  • made up of carbon, hydrogen, and oxygen atoms
  • connected in long branching chains
  • fats and oils
  • triglycerides- made up of three long fatty acid chains attached to a head group that consists of a molecule of glycerol
  • phospholipids make up the cell wall
  • not water-soluble and tend to aggregate to form droplets when placed in water
121
Q

carbohydrates

A
  • sugars
  • carbon, hydrogen, and oxygen in a ratio of 1:2:1
  • polymers made up of sugar monomers
  • glucose and fructose
  • monosccharides- carbohydrates made up of just one unit of sugar
  • polysaccharides- larger sugars, the energy storage units in both plants and animals
  • glycogen- storage of carbohydrates in animals
  • cellulose- storage carbohydrate of plants
  • can be straight chains of sugar monomers or extensively branched
122
Q

nucleic acids

A
  • contain carbon, hydrogen, oxygen, nitrogen, and phosphorus
  • polymers made up of monomers known as nucleotides
  • two major nucleic acids: DNA and RNA
123
Q

Proteins

A
  • polymers that are made up of amino acid monomers
  • 2 functional groups in common- an amino group (-NH2) and a carboxyl group (-COOH).
  • amphoteric and can act as either an acid or base
  • 20 diff amino acid monomers that make up proteins
  • polypeptides- chains of linked amino acids
  • proteins are formed by the folding of polypeptide chains
  • catalysts: all enzymes are proteins
124
Q

earth atmosphere

A

78% nitrogen, 20% oxygen, 1% argon

125
Q

the greenhouse effect

A

the buildup of carbon dioxide and other greenhouse gases in the atmosphere
- greenhouse gases are the result of the combustion of fossile fuels such as coal and oil

126
Q

acid rain

A
  • low PH

- SO3

127
Q

carbon monoxide

A
  • binds irreversibly to hemoglobin
128
Q

Conjugate pairs only differ by

A

1 H+

129
Q

diluting/concentrating a buffered solution doesn’t

A

change the PH

130
Q

equivalence point

A

enough mols of titrant to neutralize the subject acid or base, steepest part of the curve

131
Q

MV

A

MV=MV

132
Q

PH at 7 means

A

equivalence point=strong

- 7=WA

133
Q

thymol blue

A

goes red to blue at ph 7

134
Q

litmus paper

A

blue base red acid

135
Q

PKa and PKb in terms of equivalence points

A
PKa= PH (half equivalence point)
PKb= POH (half eq point)
136
Q

activated complex/transition state

A

highly unstable high-energy arrangement of atoms, must occur before products can be made and bonds can be broken, consumed in reaction

137
Q

activation energy

A

minim energy that must be supplied for the activation complex to be formed

138
Q

catalysts

A

increase rate of reaction by decreasing activation energy, aren’t consumed in reaction, and don’t change the equilibrium

139
Q

equilibrium expressions

A

solids and solvents not included

140
Q

Keq indicates

A
  • > =forward reaction favored

-

141
Q

Lechatelier

A

equilibrium

142
Q

Haber produces

A

ammonia

143
Q

only phases that can change concentration

A

g and aq

144
Q

at equilibrium RofR

A

Rates or reaction are equal but not concentration

145
Q

CA and CB how to tell

A

turn down the base

146
Q

Molarity

A

mol/liters

147
Q

Molality

A

mol/kg

- determines boiling/freezing point

148
Q

molarity and molality are both

A

measures of concentration

149
Q

solubility of gases

A
  • more pressure=more soluble

- less temp=more solubility

150
Q

Insoluble

A

Ag, Pb, Hg, OH-

151
Q

soluble

A

NO3-, ClO4-, alkali, NH+

152
Q

saturated

A

when a solvent can’t desecrate any more solvent

153
Q

kinetic molecular theory

A

KE increases proportionally when temperature is kelvin

154
Q

relationship between temperature and pressure

A

directly proportional P/T=P/T (kelvin)

155
Q

relationship between temperature and volume and law

A

directly proportional when theres no change in pressure

- Charles law V/T=V/T

156
Q

relationship between pressure and volume and law

A

inversely proportional when no change in pressure

- Boyles law P/V=P/V

157
Q

Ideal gas equation

A
  • PV=nRT
  • R=.082
  • values on the same side are inversely proportional
  • values on opposite sides are directly proportional
158
Q

intramolecular forces

A

covalent/ionic

159
Q

intermolecular forces

A

H, D-D, LDF

160
Q

bonds strongest to weakest

A

Covalent, ionic, H, DD, LDF

161
Q

what type of bond is changed in phase change

A

intermolecular

162
Q

network solids

A

covalently bonded substances that don’t consist of individual molecules (ex. diamonds and quartz)

163
Q

crystaline solids: type, bonding force, melting point, characteristics

A
  • ionic, electrostatic, high, brittle/hard
  • covalent network, shared electrons, very high, very hard
  • molecular, intermolecular, low, soft
  • metallic, sea of electrons, variable, malleable
164
Q

disposition and sublimation

A

g to s and s to g

165
Q

triple point

A

point where substance can exist as solid, gas, or liquid

166
Q

vapor pressure

A

when liquid below boiling points are evaporating

167
Q

what liquid exerts VP

A

all liquid

168
Q

atmospheric pressure

A

VP +PP of atmosphere molecules

169
Q

boiling occurs when

A

pressure above liquid is all VP or VP=atmospheric pressure

170
Q

Fluorine fun fact

A

gas at room temp

171
Q

Bromine fun fact

A

dark red liquid at room temp

172
Q

Chlorine fun fact

A

dense green gas at room temp

173
Q

Iodine

A

Solid sublimes at room temp

174
Q

Hydrogen bonds to

A

FONCl

175
Q

ideal gas temp

A

293C

176
Q

where are semimetals

A

staircase

177
Q

ionization energy

A

energy required to remove and electron

178
Q

Electronegativity

A

amount of pull an atoms nucleus has

179
Q

metallic character

A

how easily an atom loses and electron to be +

180
Q

lattice energy

A

energy required to separate a molecule of solid ionic compound into separate ions

181
Q

electrostatic force

A

force between a + and - charge

182
Q

Polyatomics

A
  • nitrate (NO3-)
  • hydroxide (OH-)
  • sulfate (SO4-2)
  • phosphate (PO4-3)
  • carbonate (CO3-2)
  • Ammonium (NH4+)
183
Q

what allows solids to conduct electricity

A

metallic bonding

184
Q

can conduct electricity in the solid state

A

graphite

185
Q

dissolution in water is highly exothermic

A

sodium hydroxide NaOH(s)

186
Q

gives off a purplish vapor as it sublimes

A

iodine

187
Q

Metalloid characteristics

A

shiny, brittle, poor electrical conductivity, and high melting point

188
Q

different electron configurations

A

DEVANIIKKK

189
Q

electron orbitals

A

probability function, quantum theory (planck), Heisenburg principle

190
Q

Bohr model

A

incorrect idea that electrons orbit the nucleus like planets orbit the sun

191
Q

Heisenburg princepal

A

electrons in orbitals, not in orbit, you can’t know the electrons position and momentum at the same time

192
Q

DeBroglies hypothesis

A

matter (including electrons) has properties of particles and waves

193
Q

outer electron location

A
  • S&P=ns & np
  • D=(n-1)d
  • F=(n-2)f
194
Q

Aufbau

A

Lowest energy state, single before double

195
Q

Pauli

A

Arrows opposite directions

196
Q

alpha decay

A

atomic # down by 2, and mass down by 4

197
Q

beta decay

A

0/-1 e, atomic mass down 1, mass number stays the same, too many protons

198
Q

gamma decay

A

table 0/0- too much energy, often accompanies other radioactive processes

199
Q

are all radioactive elements highly chemically reactive?

A

no ex. radon is a nobel gas

200
Q

do all solutes dissociate into positive and negative ions?

A

no

201
Q

rate of reaction accelerated by

A

increase of temp and catalyst

202
Q

strong acids are

A

strong electrolytes

203
Q

what does electrolysis do?

A

does not generate electricity. involves the use of electrical energy to force a chemical reaction to occur.

204
Q

Molecule

A

2+ atoms chemically bonded

205
Q

molecular formula

A

actual formula

206
Q

empirical formula

A

molecular formula simplified

207
Q

avogadros number

A

6.02x10^23=1 mol

208
Q

mass composite to empirical

A
  • find mol of each

- divide by smallest to get ratio

209
Q

entropy

A
  • measure of randomness or disorder
  • Delta S
  • -S=reaction lost entropy
  • +S=reaction gained entropy
210
Q

High S + low energy =

A

more stability

211
Q

low S + high energy =

A

less stability

212
Q

Enthalpy

A
  • energy state
  • Delta H
  • -H= exothermic
  • +H=endothermic
213
Q

Heat of formation

A

the amount of heat released/absorbed when 1 mol of a compound is formed from its elements

214
Q

Hf for elements

A

0

215
Q

Enthalpy change/ Hf for the whole reaction

A

Hf Products- Hf reactants

216
Q

Hess’s law

A

If a reaction is carried out in a series of steps, H for reactants= change in enthalpy for each individual step

217
Q

both enthalpy and entropy changes…

A

affect the value of a reactions Gibbs free energy change

218
Q

Element

A

substance that can’t be chemically broken down into a simpler substance

219
Q

atom

A

the smallest particle of an element that retains the chemical properties of the element

220
Q

nucleons

A

protons and neutrons

221
Q

CRT

A

electron gun, produces and shoots electrons through a vacuum

222
Q

JJ Thompson

A

CRT, discovered the mass and charge of an electron

223
Q

Rutherford

A

gold foil experiment

224
Q

mass number vs. atomic mass

A

mass number is the number of protons and neutrons in an atom, and it tells us about the mass of the atom in amu. Atomic mass is the average mass of all the isotopes of a certain type

225
Q

1 unit of pressure

A

1 atm=760 torr=760 mmHg

226
Q

Pressure measured with

A

barometer and nanometer

227
Q

Energy

A

The ability to do work or transfer heat

228
Q

Energy units

A

1 calorie=4.186 Joules

229
Q

Heat definition

A
  • the flow of energy from high temp to low temp

- unit: cal, J, or KJ

230
Q

Temperature

A

average Kinetic Energy

231
Q

Heat capacity

A

the amount of heat required to increase temperature by 1C

232
Q

Specific heat

A

the heat capacity of 1g of substance- how many calories to raise by 1C

233
Q

Heat formula

A

q=mc deltaT

234
Q

mass=

A

number of molecules

235
Q

may be used in combo with calorimeter to compare specific heats

A

thermometer

236
Q

completes the circuit of an electrochemical cell

A

salt bridge

237
Q

what is molality used to calculate?

A

boiling point elevation and freezing point depression

238
Q

present in liquid oxygen

A

LDF

239
Q

Responsible for the hardness of a diamond

A

network bonding

240
Q

allows copper to conduct electricity

A

metallic bonding

241
Q

colligative and example

A
  • depends only on the number of particles in a solution
  • freezing point depression
  • most particles=lowest freezing point depression
242
Q

contain undissociated aqueous particles

A

weak acid/base

243
Q

why do ionic compounds dissolve readily in water?

A

because water is a polar solvent

244
Q

c to k

A

add 273

245
Q

nonpolar and example

A
  • equal sharing

- CO2

246
Q

key element in soil

A

nitrogen

247
Q

absorbs UV rays

A

oxygen

248
Q

standard voltaic pressure

A
  • always positive for for a spontaneous chemical reaction
249
Q

zero for a crystalline solid that is elementally pure at 0K

A

entropy

250
Q

carbonates (CO3^-2) and bicarbonates (HCO3-) form

A

CO2 g when mixed with acid

251
Q

precipitation reactions occur

A
  • involve single or double displacement reactions

- between two ionic compounds

252
Q

AgCl is a

A

white precipitate

253
Q

nuclear fusion

A

responsible for energy output of the sun

254
Q

transition metal compounds are often colored because

A

the frequently possess partially filled d orbitals

255
Q

the VP of a substance is ____ of external pressure

A

independent

256
Q

isotopes

A
  • same proton, different neutron

- nearly identical chemical behavior because they have identical electron configurations (protons)

257
Q

ideal gases

A

DONT EXIST

258
Q

how is oxygen found in the bloodstream

A

as H2O

259
Q

vapor pressure of a substance depends on

A
  • substances temperature

- its mole fractions when its in solution

260
Q

which has law does temp not have to be in kelvin?

A

boyles

261
Q

positron emission

A
  • beta decay

- mass stays same

262
Q

precipitate reaction

A
  • reactants are insoluble

- not oxidation/reduction

263
Q

if equivalence point is a single equivalence point and comes at pH7

A
  • must be monoprotic

- acid must be strong

264
Q

acid/base titrations can be used to determine

A
  • alkalinity of a basic solution
  • pKa of weak acid
  • monoprotic or polyprotic
  • molecular weight of acid or base
  • NOT concentration of acid in acidic solution