Chem 20 Final Exam Flashcards
1
Q
Electro negativity
A
Used to describe the relative ability of an atom to attract a pair of bonding electrons in its valence level
Metals= low electronegativities
Nonmetals= high electronegativities
2
Q
Valence Electrons
A
Electrons occupying the highest energy level of an atom
3
Q
Valence energy level
A
The outer most energy level of an atom
4
Q
Orbital
A
A region in space around an atoms nucleus in which an electron may exist
5
Q
Valence Orbital
A
Volumes of space that can be occupied by electrons in an atoms highest energy level
First energy level: one orbital with a max of 2 electrons
Energy levels above the first have room enough for 4 orbitals (contain 0,1 or 2 electrons)
6
Q
Octet Rule
A
A maximum of 8 electrons may occupy a valence energy level
7
Q
Bonding electron
A
An atom with a valence orbital that is occupied by one single electron can share that electron with another atom
8
Q
Lone pairs
A
A full valence orbital, occupied by 2 electrons, repels electrons in nearby orbitals= 2 electrons occupying the same orbital
9
Q
Ionic bond
A
Is the attraction that results from a positive ion (metal) and a negative ion (nonmetal), 2 ions are then attracted to each other, a transfer of electrons occurs in an ionic bond
10
Q
Covalent bond
A
Is the bond that results from the electrostatic attraction between the electrons of one atom to the nucleus of another and vice versa
11
Q
Single bond
A
One pair of electrons is shared between 2 atoms
12
Q
Double bond
A
2 pairs of electrons are shared between 2 atoms
13
Q
Triple bond
A
3 pairs of electrons are shared between the 2 atoms
14
Q
Polar molecule
A
A molecule that has an overall charge separation, one end of the molecule is positive and other end of the molecule is negative
15
Q
Dipole
A
Is a result of unequal sharing of electrons in a molecule, from a difference in electronegativities
16
Q
Nonpolar molecule
A
Has no net charge separation
17
Q
Dipole-dipole forces
A
Polar molecules have dipoles, the attraction between dipoles is called dipole dipole force. The positive end of one molecule will be attracted to the negative end of a neighbouring molecule and will extend in all directions
18
Q
London Dispersion Forces
A
Result from the movement of the electrons in the molecule which generates temporary positive and negative regions in the molecule. Weak attractive forces that result when the electrons of one molecule are attracted to the positive nuclei of a nearby molecule
19
Q
Hydrogen Bonding
A
Occurs when hydrogen is bonded to a highly electronegative element (fluorine,oxygen, nitrogen). Bond is strongly polar, highly electronegative atom pulls hydrogens electron away from its nucleus, another molecules lone pair of electrons can now approach the nucleus closely on the side away from its covalent bond
*is the strongest of intermolecular bonds
20
Q
3 structures showing hydrogen bonding
A
HF(g)
Any molecule with an -OH in its structure
Any molecule with an -NH in its structure
Results in high boiling points as lots of energy must be added in order to break these bonds
21
Q
Ionic crystals
A
Crystal lattice structure
22
Q
Metallic crystals
A
Metallic bonding
23
Q
Molecular crystals
A
Molecules arranged in a regular lattice
24
Q
Network Covalent crystals
A
Network of covalent bonds
25
Physical properties of gases
1. Gases have no shape or volume of their own, they assume shape/volume of the container they are in
2. Are highly compressible (pressure increases>volume of gas decreases)
3. gases diffuse or spread out (will occupy all space that is available to them)
4. Gases mix readily with other gases
26
Pressure
A force per unit area (Pa), unit for pressure is kilopascal (kPa)
Pressure at sea level= 101kPa
1atm= 101.325kPa
27
Temperature
Is a measurement of the amount of kinetic energy (molecular motion) a sample has
28
Standard Ambient Temperature and pressure (SATP)
Laboratory conditions
P=101kPa
T=25•C or 298K
29
Standard Temperature and Pressure (STP)
P=101.325kPa
| T= 0•C or 273K
30
Boyles Law
As pressure increases, volume decreases provided that all other variables (temp, amount of gas) are kept constant
P1V1=P2V2
31
Charles Law
The higher the temperature of a substance, the greater the molecular motion of the substance. If the temp of a gas increases, the volume of the gas increases. Is valid when the pressure and amount of gas are held constant
V1/T1=V2T2
32
Absolute Zero and the Kelvin scale
Volume vs temp is graphed will result in a straight line, line will cross x-axis at -273•C (gases have no volume at this temp)
-273•C = zero kelvin = 0 K = absolute zero
Convert Celsius to Kelvin add 273, convert Kelvin to Celsius subtract 273
*always use the Kelvin unit
33
Combined gas Law
When boyles, Charles, and Gay-Lussacs laws are combined it states the relationship between pressure, Temperature, and volume of a gas
P1V1/T1=P2V2/T2
34
Law of combining volumes
States that the volume of gaseous reactants and products of chemical reactions are always in simple whole number ratios
35
Avogadros theory
Equal volumes of gas at a constant temperature and pressure contain equal numbers of molecules or atoms
SATP: 24.8L/mol
STP: 22.4L/mol
36
Ideal gas Law
Combining Charles Law, boyles Law, and avogadros theory together
pV=nRT
V=volume(L)
P=pressure(kPa)
T=temp(K)
n=#of moles (mol)
R= universal gas constant (8.314 L•kPa/mol•K)
37
Solution
A homogeneous mixture of 2 or more substances, different substances cannot be mechanically separated or seen (may be solid, liquid or gas)
38
Solvent
The substance present in the largest amount, usually determines the phase of the solution
39
Solute
The substance present in the smallest amount, what is dissolved in the solvent
40
Aqueous Solution
If the solvent is water
41
Electrolyte
A substance that conducts electricity when dissolved in water
Ex: acid, base, soluble ionic solutions
42
Non-electrolyte
A substance that does not conduct electricity when dissolved In water
Ex: molecular compounds, insoluble ionic compounds
43
Dissociation
Separations of ions from an ionic compound when it dissolved in water, molecular compounds do
Not dissociate I too ions in water
44
Ionization
Formation of charged ions from neutral atoms by the action of the solvent, acids ionize inaqueous solution
45
Concentration in Solution
```
C=n/v
C=mol/L
n=mol
V=L
Concentration is the quantity of solute divided by the quantity of solution
```
46
Concentration calculations
Always include units, make sure units cancel.
| Involve quantity of solute, quantity of solution, and concentration of solution
47
Dilution calculations
ViCi=VfCf
More concentrated to less concentrated you must add more solvent
The quantity of solute does not change, the mass and # of moles of solute remains constant after dilution
48
Dynamic Equilibrium
The rate at which particles dissolve equals the rate at which particles crystallize (2 opposing processes occur at equal rates)
49
Dissolving
When a solute dissolved in a solution, the solid particles move into the solvent to form ions
50
Crystallization
Some of the dissolved ions crystallize to become solid again
51
Saturated
When dynamic Equilibrium is reached and undissolved solute is visible even with stirring (solution will not dissolve any more particles)
Solution contains max amount of solute that can remain dissolved in a given amount of solvent at a specific temp
52
Supersaturated
Saturated solution is heated to dissolve extra solute, then cooled to original temp without extra solute crystallizing
53
Unsaturated
Solution contains less than the maximum amount of solute than can dissolve at a particular temperature
54
Factors that affect Solubility
Temperature (increase in temp will increase solubility of solids, but will decrease solubility of gases)
Pressure (increase in pressure will increase solubility of gases, has no effect on solid or liquids)
Nature of solute or solvent: solutes will dissolve better in some solvents than others, it depends on bonds between solute and solvent
55
Acids empirically defined
Taste sour, conduct electricity, react with metals to produce hydrogen gas, turn blue litmus red, pH between 0 and 7
56
Bases empirically defined
Taste bitter, conduct electricity, feel soapy/slippery, turn red litmus blue, pH between 7 and 14
57
Arrhenius Theory Acids
Acids produce hydrogen ions when they are dissolved in water, acids are substances that ionize in aqueous solution to form hydrogen ions, hydrogen ions are responsible for the acidic properties of these solutions
58
Arrhenius theory bases
Produce hydroxide ions when they are dissolved in water, are substances that dissociate to form hydroxide ions in aqueous solution, hydroxide ions are responsible for basic properties
59
Hydronium Ions
Hydrogen ion does not exist on its own but bonds to polar water molecules (H30+(aq)), is responsible for acidic properties
60
Modified Arrhenius Theory
Acidic solutions form when substances react with water to form hydronium ions, bases are substances that react with water to produce hydroxide ions
61
Acid base indicator
Substances that change color when the acidity of the solution changes, exist in 2 different forms (colours)
62
pH paper
Universal Indicator that will give the exact pH of substance (indicates acidic/basic and pH), use info from various indicators to determine pH of unknown sample or use pH meter
63
pH scale
Is a logarithmic scale that gives an indication of the nature of solution, measures relative acidity of a solution
Lower the pH the more acidic, decrease of 1 pH unit of a substance is 10x more acidic
**number of significant digits in the concentration equals the number of digits behind the decimal
64
pH calculations
pH=-log(H30+)
| (H30+)=10^-pH
65
pOH calculations
pOH= -log(OH-)
| (OH-)=10^-pOH
66
Relationship between pH and hydronium ion concentrations
There is an inverse relationship, the higher the hydronium ion concentration the stronger the acid and the lower the pH.
The higher the hydroxide ion concentration the stronger the bases and the lower pOH (low pOH=high pH)
14-pH=pOH
14-pOH=pH
67
Neutralization reaction
Double replacement reaction where an acid reacts with a base to produce water
MODIFIED: the reaction between hydronium and hydroxide ions to produce water
68
Strong acids
High conductivity, high rate of reaction with active metals and carbonates, low pH, high concentration of hydronium ions, reacts completely with water to form hydronium ions
69
Weak acids
Low conductivity, lower rate of reaction with active metals and carbonates, high pH, low concentration of hydronium ions, reacts incompletely with water to form few hydronium ions
70
Strong bases
High conductivity, high pH, high concentration of hydroxide ions, dissociate completely in solution to produce hydroxide ions
71
Weak Bases
Lower conductivity, pH closer to 7, lower concentration of hydroxide ions, react partially with water to produce few hydroxide ions, are either ionic or molecular compounds in their pure state
72
Polyprotic acids
Contain more than one acidic hydrogen ion in their formula and react more than once with water to produce hydronium ions (weak acids)
73
Monoprotic acids
Have only one acidic hydrogen ion in their formula and react only once with water to produce hydronium ions
74
Monoprotic bases
React only once with water to produce hydroxide ions
75
Polyprotic bases
React more than once with water, but all the reactions with water are much less then 50%, generally are weak bases decreases with each successive step
76
Net ionic equations
Write the chemical equation to show that entities of the dissolved ionic compounds are actually present as separate dissociated aqueous ions, leave out spectator ions (ions that are present but do not take part in the reaction)
77
Writing net ionic equations
Cancel identities that appear on both reactant and product sides, when cancelling out spectator ions they must be identical in every way- # of moles, form(atom, ion, molecule), and state of matter
78
Stoichiometric calculations
Calculations that makes use of a balanced equation
79
Gravimetric stoichiometry
The procedure for calculating the masses of reactants and products in a chemical reaction, chemical amount calculations are made from mass measurements
80
Gas stoichiometry
Extended to gases using gas volume, pressure, temperature, molar volume and ideal gas Law
Use molar volumes to convert between moles/volume if at STP or SATP(STP=22.4L/mol, SATP=24.8L/mol)
If gas is not at SATP or STP conditions then you must use ideal gas law (PV=nRT) to convert between moles and volume
81
Solution stoichiometry
The application of stoichiometric calculation principles to substances in solution, provided with the volume or amount concentration of one substance and asked to find the volume or amount concentration of another substance
82
Steps to solve stoichiometric calculations
1. Write balanced chemical equation
2. Determine the 2 substances as the known and the unknown
3. Calculate the # of moles of given substance
4. Use mole ratio to determine the unknown # of moles
5. Convert # of moles of unknown to a mass, volume, or concentration
83
Percent yield
Evaluate the overall experiment by calculating percent yield, the ratio of actual or experimental quantity of product obtained to the theoretical yield obtained from stoichiometry calculation
Percent yield=actual yield/predicted yield x100
84
Common sources of uncertainty
All measurements, purity of the grade of chemical used, washing a precipitate, qualitative measurements
85
Qualitative analysis
Identification of a specific substance present
86
Quantitative analysis
Determination of the quantity of a substance present
87
Colorimetry
Analysis by color, light emitted, absorbed or transmitted by chemical used
88
Gravimetric analysis
Uses stoichiometric calculations from the mass of a reagent
89
Titration analysis
Uses stoichiometric calculations from the volume of a reagent
90
Flame test
Some ions impart specific colours when subjected to a flame
91
Solution color
Some jones create specific colours when in solution
92
Precipitates
Using solubility table can determine what ions are in solution by whether or not a precipitate is formed
93
Precipitation reaction
A reaction where a slightly soluble product (precipitate) is formed
94
Limiting reagent
Reactant whose entities are completely consumed in a reaction, the reaction stops when and because all of this reactant is used up and none remains
95
Excess reagent
Other reactant must be present in a greater quantity than is required or you would run out of it first, some of it will remain when limiting reagent has completely reacted (to ensure it goes to completion add 10% to amount of excess reagent)
96
Determining limiting reagent
1. Write balanced equation
2. Calculate # of moles of each reactant
3. Set up mole ratio of 2 reagents
4. Find which of the 2 reagents will be used up in the reaction
5. Use molar ratio of limiting reagent and the unknown to solve
6. convert #of moles of unknown into mass/volume/concentration
97
Titration
Used to determine amount concentration of substances in solution, process of carefully measuring and controlling the addition of a solution (titrant) from a burrette into a measured volume of another solution (sample) in an Erlenmeyer flask until the reaction is judged to be complete
98
Burette
Precisely marked cylinder with a stopcock at one end, allows you to control and measure the volume of the reacting solution
99
Equivalence point
Point at which the exact theoretical amount of titrant has been added to completely react with the sample, look for a sudden change in an observable property (colour,pH,conductivity)
100
Endpoint
Point during a titration when this sudden change is observed, titration is stoped and the volume of titrant is determined
101
Observable endpoint
Use indicators that will change colour if in an acid or a base, will show momentary colour change when titrant contacts sample, closer to endpoint color change lingers, to be accurate make sure observed endpoint occurs as closer as possible to reactions equivalence point
