Chapter 4-7 Flashcards
(213 cards)
1
Q
Solution – homogeneous mixture is made up of?
A
Two or more substances
- Solute - present and smaller amounts
- Solvent - present in larger amounts
2
Q
Strong electrolyte
A
Aqueous solution- completely ionizes in solution (100% ions)
Cation & anion
3
Q
10 grams of salt in 20 grams of water
Salt water
Which is the solute and which is the solvent? ?
A
Solute- Salt
Solvent- water
Solution = salt water solution
4
Q
In aqueous solutions what is the solvent?
A
Water- (Aq) tells us water is the solvent where that something is dissolved in water
5
Q
What are three types of aqueous solutions?
A
1) strong electrolyte-completely ionizes in solution (ionization, dissociation, etc)
2) non-electrolyte
3) weak electrolyte
6
Q
Types of Strong Electrolytes
A
- Ionic compound ( metal & non-metal)
Anion & cation (positive & negative) - Strong Acid - HCl, HBr, HI, HNO3, HClO4, H2SO4 - only these (ionize fully)
- Strong Bases- group 1 Hydroxides & Ba(OH)2 from group 2 (also fully break apart in solutions- ionize)
Na(OH) K(OH) Li(OH) all strong bases
7
Q
Weak Electrolytes
A
(Partially ionize in solution) Some break into ions/ some stay togetheras molecules
- Weak Acids- all except 6 strong
- Weak Bases - anything not strong
8
Q
Non-electrolyte
A
None of the solute disassociates into ions only molecules remain
Includes everything that is not a strong or weak electrolyte
Single atoms floating around, compounds
Ar, Ne, C6H806 vitamin c
9
Q
How do u test if something is a weak, strong, or non electrolyte?
A
Using a conductivity apparatus which is a lightbulb test
Bright light= strong electrolyte (NaCl put in water gives Na+ & Cl- or
NaCl(aq) yields Na+(aq) & Cl-(aq)
Or
NaCl(s)-> H2O -> Na+(aq) & Cl-(aq)
Weak Electrolyte - glows dimly
HF(aq) double head are ->H+(aq) & F-(aq)
Double headed arrow indicates weak elec
Non-electrolyte- no light
Sugar dropped in water does break apart at all, just dissolves and shows the same molecule after arrow with aq instead of s
10
Q
What is a precipitate ?
A
The formation of an insoluable product - white solid or cloudyness formed that separates from the solution
11
Q
Soluble Componds
A
Alkali metal ions - Li+, Na+, K+, Rb+, Cs+
Ammonium ion NH4+
Nitrates - (NO3)- bicarbonates (HCO3)- & Chlorates (ClO3)- Halides Cl-, Br-, I-
Sulfates (SO4)2-
12
Q
Insoluble Exceptions - do not dissolve in water - will form a PPT Precipitate
A
Halides - silver, Mercury, lead
Ag+, Mercury (Hg2)2+, Pb2+
Sulfates that form a solid - Silver, Calcium, Strontium, Barium, Mercury, Lead
Ag+, Ca2+, Sr2+, Ba2+, (Hg2)2+, Pb2+
Halides are a binary compound formed by a halogen atom (fluorine, chlorine, bromine, iodine, or astatine) and another element or radical, such as a metal or an organic group.
Halogens:
The halogens are a group of elements in the periodic table (F, Cl, Br, I, At) known for their high electronegativity.
Binary Compounds:
Halides are compounds that contain only two different types of atoms, one being a halogen and the other being another element or a radical.
Examples:
Metal Halides: Sodium chloride (NaCl, table salt), potassium iodide (KI), and calcium fluoride (CaF2) are examples of metal halides.
13
Q
Insoluble Compounds (do not dissolve in water)
A
Carbonates (CO3)2-, Phosphates (PO4)3-, Chromates (CrO4)2-, Sulfides (S)2-, Hydroxides (OH)-
14
Q
Soluble exceptions - dissolve into ions
A
Compounds with alkali metal ions in them and the ammonia ion
15
Q
Molecular Equation
Pb(NO3)2(aq) + NaI(aq)—> PbI2 + Na(NO3)
Balance this and add states of matter
A
Pb(NO3)2(aq) + 2NaI(aq)—> PbI2(s) + 2Na(NO3)(aq)
16
Q
Ionic equation keeps solids together
A
Pb+2 + 2NO3- + 2NA+ + 2I-
—> PbI2(s) + 2Na+ + 2NO3-
2NO3-& 2Na+ are spectators
17
Q
Net Ionic
A
No spectators, only species participating in the reaction to prove a reaction took place
Pb+2 + 2I- —> PbI2(s)
18
Q
KCl(aq) + NaBr(aq)—> KBr(aq)+ NaCl(aq)
What is the ionic equation?
A
K+ + Cl- + Na+ + Br- —> K++Br-+Na++Cl-
No net bc no spectators
19
Q
Properties of acids
A
Sour taste
color changes blue lit paper red,
conduct electricity and aqueous solutions
Ions are present
react with certain metals to produce hydrogen gas (H2)
Reacts with carbonate/bicarbonate to produce CO2 gas
20
Q
Properties of bases
A
Bitter taste
Cause red lit mist paper to turn blue
Conduct electricity in aqueous solutions (makes ions Ex. OH- Hydroxide
Feel slippery
Used to make soap and detergent
21
Q
1800s Arrhenius
A
Claim that acids produce H plus ions in water and bases produce OH minus ions in water
22
Q
1932 Bronstead
A
Acids are proton donors, (ex. H+)and bases are proton acceptors (accept positively charged ions like H+)
Ex. Perchloric acid can donate a hydrogen because it has one and per Bronstead’s definition that makes it an acid because it’s a proton donor
HClO4(aq)—>H+(aq)+ClO4-(aq)
This is an example of a MONOPROTIC Acid - (yields one H+ upon ionization)
Mono = 1 protic-protons
23
Q
Diprotic Acid- yields 2H+ in 2 steps
A
H2SO4(aq)-> H+(aq) + HSO4-(aq)
HSO4-(aq)-> H+2(aq) + SO4-2(aq)
24
Q
Triprotic Acid like Phosphoric Acid will break apart hydrogen, but will require three steps one for each hydrogen
Starts with:
H3PO-4(aq)-> H+(aq) + H2PO-4(aq)
A
H3PO-4(aq)-> H+(aq) + H2PO-4(aq)
Triprotic acid yields 3 H+ in 3 steps
H2PO-4(aq)<—> H+(aq) + HPO4-3(aq)
HPO4-3(aq)<—> H+(aq) + (PO4)-2(aq)
25
Classify the following as a bronstead acid or base:
A) ClO2- B) HClO2 C)HCO3-
Bicarbonate
A) Bronstead Base (no H)
B) Chloris ACID = Bronstead Acid
C) both depending if it looses that hydrogen or gains 1 AMPHOTERIC (both)
26
Acid base neutralization
A+B yields salt water
Any acid plus any base yield salt water
Acid + base yields salt (ionic compound) +H2O(l)
27
Oxidation Reduction Reactions
Elements in elemental form by themselves have a 0 charge so Zn by itself is 0 charge but when u put it in ZnSO4 it now has a +2 charge so it oxidized
Redox - Electron Transfer reactions
SOLIDS HAVE NO CHARGE
Split reactions at bottom into 2 half reactions
Zn(s)—> Zn+2 Oxidation (adds pluses)
Cu+2—>Cu Reduction (reduces)
Zn(s) + CuSO4(aq)—> ZnSO4(aq) + Cu(s)
Ionic form
Zn(s) + Cu+2(aq) + (SO4)-2(aq)—>
—> Zn+2(aq) +(SO4)-2(aq) +Cu(s)
Net : Zn(s) + Cu+2(aq)—>Zn+2(aq) + Cu(s)
28
Oxidation Reaction - To balance half reactions, we have to add electrons bc they are electron transfer reactions
Zn(s)—> Zn+2 + 2e- to balance the charge
Zn in element form was 0 it oxidized to +2 and we added 2 electrons to balance the oxidized element
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Reduction Reaction
Cu+2–> Cu (elemental form)
How do you write the reduction reaction
Cu+2 + 2e- —> Cu(s)
We add 2 electrons to the left side of the half reaction to neutralize the Copper +2 and it yields Cu(s) in its elemental form with a 0 charge
30
Redox reaction
Zn(s) + Cu+2(aq)—> Zn+2(aq) + Cu(s)
Zn—>> Zn+2 + 2e- Oxidation
Cu+2 + 2e- —> Cu(s) Reduction
(Elemental Form)
0 Charge
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Oxidation number rules
-Element in Freeform has a zero charge
-Sum of oxidation #s equals charge if it’s an ion
-Oxidation numbers are added at the bottom to add to the total charge
-Oxygen is -2 except in H2O2 and in O2-2
-H is +1 except in LiH & NaH
-oxidation numbers can be fractional
32
Types of redox reactions
1) Combination - two elements, simple compounds form one compound
2Mg + O2 —> 2MgO
0charge+0—> +2 -2
Element form Compound
Mg —> 0 to +2 Oxidation
O—> 0 to -2 Reduction
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Type of redox reaction
Redox reaction type:
Decomposition- one compound breaks down into 2 elements/simple compounds
C—> A+B
2KClO3 —-> 2KCl + 302
34
3rd Type of Redox Reaction
Redox Reaction Type
Disproportionation: one element is simultaneously oxidized & reduced
Cl2 + 2OH- —> ClO- + Cl- + H2O
0 -2 +1 -1 -1 +1 -2
Cl= 0 to +1 is oxidation
Cl= 0 to -1 is reduction
35
4th Type of Redox Reaction is
DISPLACEMENT
Metal Displacement - any metal will displace any metal below it
Li, K, Ba, Ca, Na,……. Ag, Hg, Pt, Au
Pt= platinum
K+ NaCl—> KCl + Na
Na + KCl—> no reaction
AuCl2 + Ag —-> AgCl2 + Au
AgCl2 + Au —> no reaction
36
The common unit for concentration is?
Molarity- Capital M = moles/ liter
Moles of solute per liter of solution
37
What is the molarity of .6 moles of [NaCl] in 200 mL of solution?
BRACKETS means CONCENTRATION
Molarity=molesof solute/Liters of solution
[NaCl] in brackets means concentration of NaCl
1. Convert 200mL into Liters .200L (=200/1000=0.200L)
.6 moles of solute / .2 Liters of solution=
=3.00 moles per Liter
38
What is the molarity of .600 g of NaCl in 200 mL of solution?
Capital M
[NaCl] in brackets means concentration of NaCl
Molarity = Moles/Liters
First, we convert grams to moles
.600g NaCl*(1mol NaCl/58.5g/mol NaCl =.0103mol NaCl
200ml =.200L —> .0103mol/.2L=.0513m/L
Or .0513M or Molarity
Grams/total grams in 1 mole =# of moles
Divide #of moles by # of Liters
You get moles per liter (mol/L) = Molarity
39
Standardizing a solution
(Preparing a solution of known Molarity)
How would you create 500 mL of
1.78 M K2Cr2O7?
Convert 500 mL to .500 L
(concentration)
volume/liter X moles/liter X grams/mol
Liters requested*Mols*(#grams in 1 mole)
.500L*(1.78 mol K2Cr2O7)(294.2gK2Cr2O7
=262g K2Cr2O7
40
Steps involved in standardizing a solution
1) accurately, weigh out solute and transfer to a 500 mL volumetric flask with funnel
2) add water to dissolve the solid
3) fill exactly to the line with water
41
How would you prepare 100 mL of 0.17094 M of NaCl solution?
Standardizing a Solution
# of Liters(Molarity)(#of grams per mole)
.100L(0.17094M)(58.4gNaCl)=.998g NaCl
Add about 1g of solute [NaCl] to a 100mL volumetric flask, add a lil water to dissolve NaCl, then fill to the line with water
42
Dilution - prepare a less concentrated solution from a more concentrated one
How would you prepare 100 mL of 2.0 M H2SO4 starting with a 7.00 M stock solution.
M1V1=M2V2
7.00(?)= 100(2.0)
200/7.00=28.571 mL
Measure 28.6 mL of stock solution
Pour into a 100 mL volume metric flask
Fill exactly to the line with water
43
Loss of electrons is called
Oxidation - oxidizing, becoming positive
44
Quantitative Analysis is?
1)Gravimetric =gravity/weight/based on mass measurement
Quantitative analysis is determining the amount of concentration of a substance
Grams of precipitate formed/total grams in 1 mole of precipitate*(total grams in 1 mole of the substance who’s mass ur looking for)=grams of that substance
Take that #/grams of original sample given then multiply by 100= % mass
45
Gain of electrons is called
Reduction because they are gaining a negative charge so it’s being reduced
46
Molarity =
Moles per Liter
47
Concentration is
The amount of Solute present in a given amount of solvent or solution.
It is measured by Molarity M
M= mol/liter
(moles of solute per liter of solution)
48
Equation for Standardizing a Solution
volume in liters X moles/liter X grams/mol
=grams of solvent needed per liter of solution
Liters given X M of moles given X g/mol
49
Gravimetric -
0.5662g sample of Cl containing ionic compound is treated with excess AgNO3.
If a 1.0882g AgCl precipitate forms, what is the % mass of Cl in the original compound?
g of precipitate given/total g per mol of precipitate*(grams in 1 mol of substance ur looking for)
given g of PT/total grams*(g. in substance ur evaluating)
1.0882g AgCl/143.4g AgCl *(35.45g Cl)
=.2690g Cl Divide by original samplex100
.2690gCl/0.5662g given x 100= 47.51% Cl
50
How many grams of Cl are in 1.0882g of an AgCl precipitate?
1.0882gAgCl/143.4gAgCl*(35.45g Cl) = .2690g Cl
51
Grams of Cl calculated/ total grams of Cl in a sample X 100=
% Mass of chlorine in an original compound
gravimetric Analysis
52
What is an Acid-Base Titration?
When a standardized solution is gradually added to another solution of unknown concentration until the chemical reaction between the 2 solutions is complete
53
Acid-Base Titration
Buret (long glass tube) is filled with an acid or base of unknown concentration.
The opposite substance (acid or base of known concentration) is put into a flask set below the buret.
Open the valve of the buret to allow some solution into the flask set on top of a magnetic plate, to react with the opposite substance while stirring the mixture
With an indicator like phenopthalein to tell us when the acid and base have reacted with each other by turning pink (neutralizing the base completely in the flask)
54
A student finds that 0.7339 g KHC8H4O4 is needed to neutralize 30.25mL of NaOH. Find the concentration of NaOH.
Grams given /total g per mol of given *(1mol KHP/1mol NaOh) =moles of NaOH
Then: moles/volume=Molarity
0.7339g KHP/204.2g KHP *(1 mol NaOH/1mol KHP {ratio}) = .003594 mol NaOH
.003594 NaOH/.03025L=.1188M
Molarity=moles/liters
55
What is it called when any substance is heated to a high enough temperature (or added energy) that it glows
Electromagnetic radiation
ER
56
Who studied and found that light travels thru space in ___?
Waves
1873 James Maxwell
57
Distance between 2 adjacent peaks in a wave is called its
Wavelength
58
Wavelength is measured in?
Nanometers
59
1 nano=
10^-9 meters
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# of crests or cycles that pass a point per second is called
Frequency (how frequent a crest passes a point per second)
61
Frequency is measured in?
Hertz (hz) s^-1 1/s (#of crests per second)
(Bird’s beak)
More small waves = high frequency
Big waves, less pass= lower frequency
62
The unit for frequency is?
Greek symbol for Nu
Looks like (birds beak)
63
Wavelength is measured in?
Greek letter Lambda (looks like teepee)
64
Frequency measured
Per seconds or 1/s or s^-1
Or hertz or hz
Greek letter nu (birds beak)
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Vertical distance from the midline to the peak or trough of a wave
Amplitude
66
Total amplitude from the top of the wave to the very bottom =
2A
A=amplitude= half the height of the wave so the full height is 2A or 2 Amps
67
Light waves travel at a given speed with the symbol
U
And u = lambda X Nu
Speed = distance/time
= (distance/wave)*(wave/time)
=wavelength*frequency = u (speed/wave)
68
Speed of light =
C = 3.0x10^8 m/s
69
Colors of the spectrum
Roy G Biv
Red, orange, yellow, green, blue, indigo, violet
70
Electromagnetic Spectrum ?
Visible Light=
Gamma Rays, X-Rays, UV Rays, Visible, Infrared, Microwaves, TV/Radio Waves
GXUV IMT
Visible = 400-700nm
71
What can be used to separate the colors of visible light?
Prism
72
RoyGBiv goes from what frequency ?
Low to high frequency from red to violet
Low frequency- LONG (longer wavelength)
High frequency- SHORT (Short wavelength
73
How are frequency and wavelength related?
They are inverses of each other
74
RoyGBiv is located in which part of the spectrum?
Visible light 400-700nm but it’s backwards vib G yoR
75
Wavelength on the spectrum
Gamma=short/high energy
to Tv/Radio=long/low energy
High energy can do major damage (can destroy bombs)
Microwaves =molecules spin causing friction that produces heat
76
Infrared means
Beyond red
77
Ultra violet means
Beyond violet
78
Where is violet on the visible spectrum?
400 nm
79
Where is Red on the visible spectrum?
700 nm
80
Classical
What everyone believes to be true up to a certain point
81
Classical physics
Assumed atoms and molecules could emit/absorb any amount of energy
82
1900 Max Planck
Atoms/molecules can absorb/emit energy in specific amounts called quantum
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Energy of 1 quantum of energy
E=planks constant*frequency
E =h*nu
E(energy)
h (plank’s constant = 6.63 x 10^-34 Jules*seconds)
Nu (birds beak) = frequency
Energy= 6.63x10^-34*(frequency)
84
Speed (c) =
(Lamda)*(frequency)
Teepee*birds beak (nu)
85
H (plank’s constant) =
C (speed of light)/lambda (wavelength)
H=C/lambda
86
C=
Speed of light
87
Energy =
h*c / lambda
(Plank’s constant)*(speed of light) /wavelength
Energy = h*nu
Nu=speed of light/lambda
88
Photoelectric Effect -
Einstein 1905- Electrons are ejected from the surface of certain metals exposed to a light of at least a minimum frequency (called the threshold frequency)
-Light continuously shines down and electrons are continuously ejected from the metal
-Brighter the light, more electrons are emitted
-Beam of light acts like a steam of photons
89
Wave-particle duality
Light can act like a WAVE but can also act like a PARTICLE
90
Examples of photoelectric effect
TV Remote (IR Light hits electronic TV eye)
Motion sensors (beam of light)
Automatic hand dryers (use light beam)
91
Continuous vs. line spectra
Emission spectra - when a substance is heated, a prism separates light into its component colors
92
When a solid is heated, like a piece of iron or another element, we will see?
A continuous spectrum like a rainbow (colors blended together continuous)
93
An element or sample in a gas state would create a
Line spectra or discrete spectra
( with different colored ones at distinct spots that can help identify elements by their line spectra)
94
We examine an emission spectra with a
Spectrometer or spectrophotometer
95
A spectrometer is made up of
Eyepiece
Prism
Ruler
96
Goal of Bohr’s Model
To explain how the electron in H was responsible for energy emitted in the form of light
97
Two assumptions of Bohr’s model
1. The electron revolves around the nucleus and stable circular orbits.
2. Energy of the H atom is quantized and travels in one of many different orbits called energy levels designated by the “principal quantum number,” n
98
Energy levels, which require a certain amount of energy that the electron must have to stay in one of these levels, can be calculated with the formula
En= -Rh(1/n^2)
Energy Levels
N=1 lowest (ground level)
N=2 1st excited state
N=3 2nd excited state
Rydberg Constant: Rh=2.18*10^-18J
E sub n = negative rh*1/n-squared
99
What to plug in for n squared in the energy formula
1) Plug in 1 for n if
2) Plug in 2 for n if
3) Plug in n=3 if
1) 1 if the electron is in the lowest energy level
2) 2 if the electron is in the 2nd level (1st excited energy state)
3) 3 The electron is in the 2nd excited energy state
100
What is the value for R and what is it called in the energy formula
R= Rydberg Constant Rh=2.18*10^-18J
101
To calculate the energy of a transition
The Change in E=
Delta🔺E= Ei-Ef
Ef= -Rh(1/nf^2) 1/n-final squared
Ei= -Rh(1/ni^2) 1/n-initial squared
🔺E= Rh(1/ni^2-1/nf^2)
The change in Energy= Rydberg Constant*(1/n initial squared -1/n final squared)
102
Electron sits in n=1 regularly, we add energy and it jumps up to n=2 and stays there as long as it has enough energy to maintain it.
-as soon as the electron looses a little bit of energy, it drops back to the ground state at n=1 & the extra energy it looses from dropping is given off as light
This light is related to the wavelength which is related to the color that’s given off.
This is a good example of how fireworks work. The fireworks give off light as they drop down. So different elements are used to give off the different colors
103
Scientists who studied the wavelength of light that corresponds to the transition of the electron
Name each scientist & the n-final for each
LBP-BP - know the n final for each!!!!
Lyman n=1 to n=5 (n final is 1)
Balmer n=2 to n=5 (n final is 2)
Paschen n=3 to n=5 (n final is 3)
Bracket n=4 to n=5 (n final is 4)
Pfund n=5
Solid black line at the top is a band called a continuum
104
How do u know if energy is emitted or absorbed?
When the sign of 🔺E is negative, energy is released (negative=loss of energy)
Or if the electron is dropping from a higher level to a lower level, it is giving off energy in the form of light so light is emitted
105
Would u be able to see light being emitted?
You would have to calculate and if the calculation comes to 400-700nm, that’s the only visible light you can see
106
Difference between E and Delta E
E is the absolute value of Delta E
Delta tells you the sign, E tells you how much was taken out, just a number, not a positive or negative sign
107
Who concluded from experiments that waves can be behaved like particles and particles can have wavelike properties in 1924
Louie de Broglie
He plucked guitar strings and noticed the waves and how it was related to particles
His formula was lambda=h/mv
M=mass in kg
V=velocity (m/s)
108
Calculate the wavelength of an electron moving at 62 m/s
We can use Louie de Broglie’s formula
Lambda= plank’s constant/mass of an electron in kgs*(62m/s)
Lambda= 6.63x10^-34J/[9.1x10^-31kg*(62m/s)]= 1.2x10^-5
Divide this by 10^-9 to convert to nm
=1.2x10^4nm =infrared region, not visible
109
Calculate the wavelength of a tennis ball moving at 62 m/s
THIS EQUATION NO LONGER WORKS FOR LARGER ITEMS LIKE THIS
Lambda= h / [(mass of tennis ball)(speed)] multiplied by conversion factor 1kg/s^2/J
Lambda= (6.63x10^-34J/s)/[(6.02x10^-2)(62m/s)] X (1kg/s^2/1J)= 1.8x10^-25nm this is no where on the spectrum
110
Who shot electrons through gold foil and found that electrons have wave-like properties from an x-ray like pattern?
Davisson & Germer
111
Momentum =
Mass * Velocity
112
What principal claimed that it’s impossible to know both momentum and position of a particle with certainty
The Heisenberg Uncertainty Principle
113
A particle can move you can determine both its momentum and its position, but not both with certainty
Who claims this?
Heisenberg Uncertainty Principle
114
Studies the probability of finding an electron in a certain region
Schrödinger with Quantum Mechanics
115
No two electrons and an atom can have the same set of four quantum numbers is what principle?
The Pauli exclusion principle
116
Each electron on the periodic table is gonna have a unique set of four numbers and this is the
Pauli Exclusion Principle
(Similar to each person having a unique home address)
117
Given the symbol n and gives the shell or orbital, and tells the average distance of the electron from the nucleus
Principal quantum number
N=1 closest to nucleus
N=2 is further from the nucleus etc
Goes to infinity orbitals
118
This number tells the shape of the orbital and gives the subshell
L = n-1
Angular momentum, quantum number, L
Value Ranges from 0 to n-1
So if n = 1, L has to be 0
N=2, L = 0 or 1
N=3, L = 0, 1, or 2 (these are boxes on the rings that are occupied by electrons)
119
Tell special orientation, gives available energy levels within a subshell
Magnetic quantum #, mL (m sub L)
- ranges from negative to positive l
L=0, mL=0
L=1, mL= -1, 0, 1
L=2, mL=-2,-1,0,1,2
120
Last quantum number which determines electrons have two different spins
Electron spin, quantum number, M sub S
Stern & Gerlach (1924)
Two different spins- clockwise & counter clockwise +1/2, -1/2
121
States that spinning charges generates magnetic fields
Electromagnetic theory
122
Quantum # chart
123
Predict the number of subshells in the fourth shell for n=4
4s, 4p, 4d, 4f
4 different subshells for n=4
124
How many orbitals are in each of the 4s, 4p, 4d, & 4f subshells ?
4s = 1 orbital
4p= 3 orbitals
4d = 5 orbitals
4f = 7 orbitals
= 16 total orbitals
Each box holds 2 electrons so there could be 32 electrons
125
Give all possible subshells and designations when n=3
When n=3, L= 0, 1, or 2
3s, 3p, 3d
126
How many orbitals and electrons can have the following sets of quantum numbers?
A) n=2
B)n=2, l=0
C)n=2, l=2
D)n=10
A) n=2 (2s, 2p, 2p, 2p) = 4 orbitals 8 electrons
B) n=2, l=0 (this is a 2 s orbital with 2 electrons
C) n=2, l=2 (l cannot be 2, only 0 to n-2)
D) n=10 (square it = 100 orbitals) electrons = 200
127
how many electrons can have these quantum numbers?
N=2, l=0, ml=0, ms=-1/2
1 electron according to Pauli
128
List all the quantum numbers with the designation 3d
N= 3 bc it’s 3d
L= 2
Ml can be -2, -1, 0, 1, or 2
Ms= +1/2, -1/2
129
List all of the sets of quantum numbers for 1s
1 orbital 2 electrons means 2 sets of quantum #s
1, 0, 0, +1/2
1, 0, 0, -1/2
130
# of Orbitals is
1/2 the # of elements in that subshell
131
M sub l values range from
Negative to positive l
So if l is 2, m sub l is -2, -1, 0, 1, 2
132
How many electrons can fit in an orbital?
2
so it’s safe to consider an orbital as 1 of those boxes with the opposite facing arrow spins
133
What does an orbital look like when drawing a diagram?
A box with arrows in it that represent electrons and their spins as +1/2 or -1/2
134
What is a subshell designation ?
The subshells are the designations labeled 1s or 2p or 3d or 4f, etc
135
Total number of orbitals =
1/2 the # of elements in that shell
So if there are 6 elements, that means there 3 orbitals
2 electrons can fit in each orbital
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N=
The row # on the periodic table
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How many electrons can fit into an orbital?
2
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How many orbitals are in 2p?
3d?
2p = 3 orbitals (6 elements)
3d = 5 orbitals (10 elements)
# of Orbitals is 1/2 the # of elements in the subshell
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How many electrons are in an orbital?
How many electrons are in the 1s orbital?
The # of elements in the orbital happens to correspond to the # of electrons in that orbital
2 electrons in the 1s orbital.
1- Hydrogen
2- Helium
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How many ORBITALS are in n=3
18 elements in 3rd row, 9 ORBITALS
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How many electrons are in n=2?
8 elements, 8 electrons
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How many electrons are in N=3 ?
18 elements 18 electrons
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What shape is the s orbital?
Sphere
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What shape are the 3p orbitals?
Dumbbell shaped and sit on the x-axis horizontally, the z-axis vertically, and the y-axis which pokes thru the computer screen diagonally towards the right upper back and left lower front (looks like a pair of lungs if you ask me)
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What shows the distribution of electrons in atomic orbitals?
Electron configurations
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What tells us which orbitals we can find electrons?
Quantum numbers
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How the probability of finding an electron in s, p, or d orbitals would look
Looks like 4 eggs standing up long way in a circle next to each other, all sides touching and the z-axis running north and south out of the center of the circle, the y and z axis like an x shape centered over the z
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What states that the most stable arrangement of electrons is the one with the most parallel spins
Hund’s Rule
Arrows placed in the orbitals facing the same way bc there is only 1 per box until all boxes are filled and the 2nd round is added to each box
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Electrons fill lowest energy orbitals first - is a rule of what?
Aufbau Principle
Aufbau-German word means”building up”
Start at bottom and work your way up. Start at bottom, fill lowest energy orbitals first, when they fill, then move up to next
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Which energy level is closest to the nucleus?
N=1
The values for n increase as the distance from the nucleus increases and the energy it takes to travel higher increases
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Spaces between energy levels (increases or decreases) as you travel further from the nucleus to higher n=levels?
Spaces between energy levels decreases, becomes smaller, as the energy levels get further from the nucleus to larger n=levels.
A 3D energy level is more energy than a 4S level and 4p is more than both of them
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Electron Configurations
Hydrogen
Helium
Lithium
Boron
Neon
Argon 18 electrons
Zinc
H 1s1 = 1s shell #1
He 1s2 = 1s shell #2
Li 2s1= 2s shell#1 or [He] 2s1 Nobel gasSC
B 1s2,2s2,2p1 or [He] 2s2,2p1
Ne 1s2,2s2,2p6
Ar 1s2,2s2,2p6,3s2,3p6 or [Ne] 3s2,3p6
Zn 1s2,2s2,2p6,3s2,3p6,4s2,3d10
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Where do electrons go first? 3D or 4S?
4s bc it’s a lower energy level than 3d
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Noble Gas Shortcut of Electron Configurations
Previous noble gas in brackets followed by the electron configuration of the row leading up to the element you’re figuring.
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Exceptions bc:
HALF FILLED & COMPLETELY FILLED
D-Orbitals are very stable
So it will take 1 electron from the previous shell to make its d shell full or 1/2 full to d5 or d10 to be stable if the d shell is 1 short of being 1/2 full or completely full
We expect Cr to be [Ar] 4s2,3d4 but what we observe is [Ar] 4s1,3d5
Mo & W in D orbital behave the same way
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Cu what u expect? what it is ?
Ag what u expect? What it is?
Also Au for gold does this
Expect Cu = [Ar] 4s2,3d9
Observe Cu = [Ar] 4s1,3d10
Expect Ag = [kr] 5s2,4d9
Observe Ag = [kr] 5s1,4d10
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A pictorial representation of electron configurations with boxes representing orbitals and arrows representing electrons
Orbital Diagram
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Hydrogens orbital diagram
H 1s1
A box labeled 1s under the box and 1 arrow (pointing EITHER upwards OR downwards) inside the box (to represent its’ electron spin)
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Pictorial Diagram of Helium
He 1s2
A box with 1s labeled under the box and 2 arrows inside the box (1 arrow pointing upwards and a 2nd arrow pointing downwards) to represent 2 electrons with opposite spins
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Phosphorus pictorial diagram
1s2 2s2 2p6 3s2 3p3
1s box with arrow up & arrow down
2s box with arrow up & arrow down
2p- 3 boxes connected with 2 opposite
pointing arrows in each box
3s box with arrow up & arrow down
3p - 3 boxes connected, with 1 upward
facing arrow in each
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Substances attracted by a magnet (their electrons are unpaired)
All ODD atomic #’s ATTRACT to even out-
PARAMAGNETIC - odd # of electrons
If unpaired, they want to attract a partner
Unpaired (para-1) attracted-for a magnet- para- magnet .. for-a-magnet
As long as 1 unpaired electron - paramagnetic
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Substance is repelled by a magnet (electrons are all paired)
EVEN atomic #s are NOT ALWAYS
DIAMAGNETIC - ONLY if they are all paired, they repel others bc they are happily paired
they can have an even # that’s spread across orbitals unpaired to fill the orbital b4 pairing up so look at last orbital drawn
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19th century scientists arranged elements by:
Atomic mass
(Mass related to chemical behavior)
Example - density is a higher value for a heavier mass, boiling point is higher also
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1864 - every 8th element has similar properties. What is this called and who was the scientist?
John Newlands “Law of Octaves” (not really a law bc it doesn’t always work)
P columns- the element below each 1
& 3s1/4s1 Na/K & 3s2/4s2 Mg/Ca
Lithium & Sodium - similar properties
Beryllium & Magnesium = similar properties
Boron & Aluminum
Carbon & Silicon
Nitrogen & Phosphorus
Oxygen & Sulfur
Neon & Argon
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Who improved Newland’s “Law”?
5 years later 1869 Dimitri Mendeleev
-grouped 66 elements by property
-left spaces for future elements
-predicted unknown future elements properties
Ex. Predicted Eco-aluminum/4 years later it was discovered as Gallium 69AMU predict & its 69.723
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A year after Mendeleev, what scientist came along and created a similar table?
1870 Lothar Meyer
167
What element honors Mendeleev?
Mendeleevium
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1913 Henry Moseley
Unique X-Ray pattern for each element - and that’s how each element got its Atomic #
He Discovered a relationship between the atomic number and the frequency of x-rays generated when elements bombarded with electrons
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Representative elements of MAIN GROUPS on the periodic table are
1A-7A
Each has an incomplete either S or P outer shell
170
Which group has completely filled p orbitals except for Helium
Noble Gases
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Electron configurations and group numbers
The last set of configuration for an element tells you its group number
- or the # of electrons in its OUTER shell is its group number
Li 2s1 group 1
Na 3s1 group 1
Al ends in 3p1 in group 3A
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Elements that have similar properties also have
The same number of valence electrons
(Electrons is the valence (outer most) shell
173
If we form a cation, we
Loose electrons so there’s 1 less electron
174
If we form an anion,
We gain electrons so there’s 1 more electron
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If an element gains an electron, it becomes the element to the ___of it.
If it looses and electron it becomes the element to the _____ of it.
If an element gains an electron, it becomes the element to the RIGHT of it
If an element loses an electron, it becomes the element to the LEFT of it
Flourine becomes an ion by gaining a -1 charge to become F- = Neon (noble gas)
Na looses an electron to become Na+ to behave like Neon to become noble
176
Elements that loose or gain electrons to behave like the element noble to them become
Isoelectronic with each other
F- & Ne are isoelectronic with the same configurations
Na+ & Ne are isoelectronic
Must use the plus and minus sign in these
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Loose electron
Cation
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Gain electron
Anion
Gains a negative
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When we are looking at cations for transition metals, how are electrons removed
First from s orbital, then from d orbital
Ex.
Mn [Ar] 4s23d5
becomes Mn+1 [Ar] 4s13d5
If we take a 2nd electron it comes from the 1st (outer most shell) Mn+2 [Ar] 3d5
Then they come from the d shell
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What is 1/2 the distance between two nuclei in two adjacent atoms
Atomic Radius -
Which is determined by:
ATTRACTION of NUCLEI and VALENCE Electrons TO the PROTONS in the Nucleus
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Strong Attraction between atoms means?
Weaker attraction means?
STRONG attraction leads to SMALL radius
WEAK attraction leads to LARGER radius (not held together as tightly)
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Atomic Radius of elements on the periodic table
Grow left and downwards
So the bottom left elements have larger atomic radius than things on the top right
183
What effect a physical and chemical properties of ionic compounds
Ionic radius- radius of an anion or cation
If the atom forms a negative anion, the radius gets larger as it repels apart
If the atom forms a cation, we are taking away electrons from the outer shell, less repulsion - smaller radius
184
Arrange elements in order by increasing radius
The radius grows left and down according to arrangement on periodic table
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If we want to compare Nitrogen-3 & F-
First look at the number of protons and electrons
-Nitrogen has 7 protons with -3 charge so 3 extra electrons makes 10 electrons
-Flourine has 9 protons with -1 charge so 1 extra electrons makes 10 electrons
Nitrogen -3 is larger bc less pull with less protons to electron charged
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We we have equal number of protons but electrons are different. What do we look at to determine the atomic radius?
Repulsion
The outer shell that has more electrons has more repulsion pushing away from each other making them larger
187
What is the 3rd ion property
Ionization energy - minimum amount of energy needed to remove an electron from a gas atom in its ground state
X—> x+1 + e-
X yields x with a +1 charge, plus a negative electron - this is 1st ionization energy
188
If we take a 2nd electron away from x that became x+1
We start out with X+1(g) —> x+2(g) + e-
X+2 means removed 2 electrons and called the 2nd ionization energy
189
The 100th ionization energy would be
X+99(g) —> X+100(g) + e-
= 100th ionization energy
190
What is the 5th ionization energy of phosphorus?
P(g)+4 —> P(g)+5 + e-
Gains electrons on right of arrow is ionization energy
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The trend for ionization energy is opposite of atomic radius
It increases up and to the right
It increases right bc it gets closer to a full shell becoming a noble gas, harder to take away an electron
Closer to left, easier to loose electrons to move Back to previous noble gas
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Last periodic property is
Electron Affinity - the energy charge associated by the acceptance of an electron in the Gas state
193
Ionization energy _______ electrons,
Electron Affinity __________ electrons
Ionization energy LOOSES electrons,
Electron affinity GAINS electrons
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If we take some element X in the gas state, we add an electron to it, it becomes?
Which is an example of ?
X(g) + e- ——> X-1(g)
Looses electrons left of arrow = electron affinity
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Add a 2nd electron
X-1(g) + e- —> X-2(g)
X-99(g) + e- yields X-100(g)
Electron affinity increases up to right
196
High electron affinity means
Wants to gain electrons
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Attraction between electron and proton is greater
At the top of the periodic table, the more levels we fall down, the less attraction bc the electrons are so many levels out away from the protons in the nucleus
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Li 1s2, 2s1 & Be 1s2, 2s2
Which has the highest
1) 1st ionization energy
2) 2nd IE
3) 3rd IE
The first ionization energy is the 1 that would require MORE energy to remove the outermost electron from a neutral atom
To assess which element has the highest third ionization energy between lithium (Li) and beryllium (Be), we need to consider the electron configurations and the effects of removing electrons:
• Lithium (Li):
• Electron configuration: 1s² 2s¹.
• After the first ionization, it becomes Li⁺ (1s²).
• After the second ionization, it becomes Li²⁺ (the 1s shell is fully occupied).
• The third ionization involves removing an electron from the 1s² configuration, which requires a significant amount of energy due to stability.
• Beryllium (Be):
• Electron configuration: 1s² 2s².
• After the first ionization, it becomes Be⁺ (1s² 2s¹).
• After the second ionization, it becomes Be²⁺ (1s²), similar to lithium’s stable configuration.
• The third ionization would involve removing an electron from the 1s² configuration as well, requiring a high amount of energy.
Both elements will have high third ionization energies due to the removal of an electron from a full 1s subshell. However, the third ionization energy for lithium is expected to be higher because it is removing an electron from a more stable configuration (1s²) after the first two ionizations, while beryllium’s second ionization leaves it with a single 2s electron before the third ionization.
Thus, lithium (Li) has the highest third ionization energy compared to beryllium (Be).
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Given Li, Be, B, O, F, Ne, which has the:
1) highest first electron affinity
2) highest second electron affinity
3) lowest third electron affinity
4) lowest fourth electron affinity
Based on the elements provided (Li, Be, B, O, F, Ne), let’s evaluate each option:
1. Highest 1st EA (Electron Affinity):
• Fluorine (F) typically has the highest electron affinity among these elements because it is highly electronegative and readily accepts an electron.
2. Highest 2nd EA:
• Oxygen (O) generally has a higher second electron affinity than the others, as removing an electron from a negatively charged ion (O⁻) requires more energy than from a neutral atom.
3. Lowest 3d EA:
• This term seems to be a bit unclear, but if it refers to the energy required to add an electron, then Boron (B) may have the lowest because it has fewer electrons and a less stable configuration compared to the others.
4. Lowest 4th EA:
• Neon (Ne) would likely have the lowest fourth electron affinity since adding an electron to a noble gas configuration is energetically unfavorable.
Summary of Answers:
1. Highest 1st EA: F (Fluorine)
2. Highest 2nd EA: O (Oxygen)
3. Lowest 3d EA: B (Boron)
4. Lowest 4th EA: Ne (Neon) {should have been nitrogen but it wasn’t an option }
BORON is the lowest 6th electron affinity
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What is a halide?
Halides are a binary compound formed by a halogen atom (fluorine, chlorine, bromine, iodine, or astatine) and another element or radical, such as a metal or an organic group.
Halogens:
The halogens are a group of elements in the periodic table (F, Cl, Br, I, At) known for their high electronegativity.
Binary Compounds:
Halides are compounds that contain only two different types of atoms, one being a halogen and the other being another element or a radical.
Examples:
Metal Halides: Sodium chloride (NaCl, table salt), potassium iodide (KI), and calcium fluoride (CaF2) are examples of metal halides.
201
%mass
Grams/total grams X100 = % mass
202
What type of bond forms when Adams share electrons between nonmetal atoms?
Covalent Bonds:
These bonds form when atoms share electrons to achieve a stable electron configuration, often occurring between nonmetal atoms.
203
What type of bond forms when two atoms share one pair of electrons?
(2 electrons total)
Single Bond:
Two atoms share one pair of electrons (two electrons).
Example: Hydrogen gas (H₂).
204
What type of bond forms when two atoms share two pairs of electrons?
(4 total electrons)
Double Bond:
Two atoms share two pairs of electrons (four electrons).
Example: Oxygen gas (O₂).
205
What type of bond is formed when two atoms share three pairs of electrons?
(6 total electrons)
Triple Bond:
Two atoms share three pairs of electrons (six electrons).
Example: Nitrogen gas (N₂).
206
What is the difference between the strength and length of bonds?
Triple bonds are stronger and shorter than double bonds, which are in turn stronger and shorter than single bonds.
207
Sigma and Pi Bonds
A single bond is a sigma bond.
A double bond consists of one sigma bond and one pi bond.
A triple bond consists of one sigma bond and two pi bonds.
208
Examples of Molecules with Different Bond Types:
Single Bonds: H₂O, Cl₂, CH₄.
Double Bonds: O₂, CO₂, C₂H₄.
Triple Bonds: N₂, C₂H₂.
209
Standardizing a solution
The liters given X Molarity given X (the total grams in 1 mole of given substance)
210
Lambda x Nu =
U (the speed of light)
Wavelength X speed
211
Lamda=
Distance/wave (wavelength)
212
Nu=
Wave/time (speed)
213
U=
The speed of light
Lambda x Nu
Wavelength X Speed (wave/time)
