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exam 3 Flashcards

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1
Q

Write the full electron configuration for aluminum (Al)

A

1s^2s^2 3s^23p^1

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2
Q

How many core and valence electrons does aluminum have?

A

10 core 3 valence

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3
Q

What is ionization energy?

A

the energy required to remove an electron from an atom in the gas phase

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4
Q

Is it easier to remove a core or valence electron from an atom?

A

It is easier to remove a valence electron because a valence electron is further from the nucleus and has a weaker force of attraction to the nucleus. A weaker force of attraction needs LESS energy to overcome the force (lower ionization energy). A core electron is closer to the nucleus and has a stronger force of attraction to the nucleus. A stronger force of attraction needs more energy to overcome the force (higher ionization energy).

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5
Q

Draw a graph showing the first six ionization energies of Aluminum. Explain how this graph represents aluminum (explain changes in force and energy)

A
  • You would draw the first 3 dots being closer to the x-axis, and the 4 dot in the middle of the graph and the other two linearly increasing until 6.
  • There is a big jump in ionization energy between the 3rd and fourth ionizations. Aluminum has three valence electrons. ve- are further away from the nucleus and experience a lower effective nuclear charge than core electrons, making them less attracted to the nucleus and easier to remove with less energy. The fourth ionization requires more energy because of the attraction between the core electrons and the nucleus.
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6
Q

draw a PE curve for h atom and then he atoms. Identify the curve and label the axes.

Explain the depths of the potential wells.

Explain the positions.

A

draw a big dig closer to the PE line for H and a small dip closer to the r (or the right side) and label that He

the pe well is deeper for H because when two hydrogen atoms interact they form a covalent bond, while 2 He atoms interact through LDFs. Covalent bonds are much stronger than LDFs.

The potential well for He has a longer distance between the nuclei than H because when two atoms are interacting through LDFs they are further apart then when they are covalently bonded together.

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7
Q

The melting point for Iodine (I2) is 387K and the boiling point is 457K. Draw an atomic/molecular level representations of Iodine at 400K and 460K. Within each drawing indicate (with an arrow) and label any bonds and/or interactions present. If no bonds and/or interactions are present.

A

I2 at 400K (liquid) - draw ovals representing dipoles with two iodine molecules bonded together. (dont forget the d signs) Draw a arrow inside

12 at 460K (gas) - show two ovals separated and a covalent bond still exisiting.

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8
Q

The melting point of Xe is 161K and the boiling point is 165K. Draw Xe at 163K and 170K.

A

Xe at 163K (liquid) - only LDF is present. small circles of solo bonds.
Xe at 170K (gas) - (none) just two circles not touching.

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9
Q

What bonds/interactions are overcome when iodine (i2) and xenon (Xe) boils?

A

LDF

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10
Q

Iodine has a higher boiling point than Xe. What is the evidence and reasoning?

A

evidence:
The electron cloud for a Xe atom has 54 electrons. The electron cloud for an iodine molecule (I2) has 106 electrons.
reasoning:
More electrons in the electron cloud result in bigger dipoles (a bigger 𝛿+ and 𝛿−). Bigger dipoles result in a stronger force of attraction (stronger LDF) - remember Coulomb’s law. Stronger LDFs need more energy (a higher boiling point) to break.

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11
Q

a. List the elemental forms that have relatively low melting and boiling points? Identify if the element is a metal or a nonmetal

b. What type of bonding and/or interactions might be overcome during a phase change for each of the elemental forms you listed in part a?

A

H2, He, N2, O2, F2, Ne (nonmetals)
- LDF

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12
Q

List the elemental forms that have relatively high melting and boiling points? Identify if the element is a metal or a nonmetal.

What type of bonding and/or interactions might be overcome during a phase change for each of the elemental forms you listed in part c?

A

Li, Be - metal
B – nonmetal (metalloid)
C - nonmetal

-covalent bond or metallic bond

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13
Q

What elements overcome LDF during a phase change?

A

H2 (small mol) , He (discrete atoms) N2 (small mol), O2 (sm mol) F2 (sm mol), Ne (discrete atoms)

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14
Q

What elements overcome Covalent Bonds during a phase change?

A

B(s) , C(s) (extended network)

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15
Q

What elements overcome metallic bonds during a phase change?

A

Li(s) , Be (s) (extended network)

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16
Q

What pattern do you see regarding the melting and boiling points of these elements relative to the types of bonding and interactions overcome during a phase change?

A

Elemental forms with very high melting and boiling points are extended network solids. The atoms are held together by metallic bonding or covalent bonds.

Elemental forms with lower melting and boiling points are made of either discrete atoms or small molecules with LDFs between them.

The extended network solids have higher melting and boiling points because both metallic bonds and covalent bonds are very strong interactions. These strong interactions must be broken when the solid melts (or the liquid boils). It would require a lot of energy (high temperature) to break these interactions.

LDFs are relatively weak interactions. Less energy is required (lower temperatures) to overcome these interactions. It is the LDFs (not the bonds within the molecules) that are overcome when the solid melts (or the liquid boils).
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17
Q

Why is nitrogen a gas at room temperature and carbon is a solid?

A

The nitrogen exists as small, non-polar diatomic molecules interact with each other through LDFs caused by the momentary fluctuating dipoles (instantaneous diploes induce dipoles in nearby molecules – weak). Carbon atoms form localized, strong interactions between the atoms (extended networks in a lattice structure - think diamond or graphite). At room temperature there is enough energy to break the LDFs between the nitrogen molecules but there is not enough energy to break the stronger bonds between the carbon atoms.

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18
Q

Draw a molecular level picture of Copper and use it to describe the bonding present

explain how the model of bonding you discussed explains the properties you listed for each substance

A

Copper - high MP, has a color, malleable and ductile, shiny, and conducts electricity

Lattice of regularly spaced nuclei and core electrons. Valence electrons are delocalized and can move freely because they are in molecular orbitals that span the system.
you would draw about 12 spaced atoms with positive charges drawn in the middle and then delocalized valence electrons.

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19
Q

Draw a molecular level picture of Graphite and use it to describe the bonding present

explain how the model of bonding you discussed explains the properties you listed for each substance

A

Graphite - high MP, soft, slippery, shiny, conducts electricity.

Two-dimensional sheets of sp2 hybridized C atoms. The remaining p-orbitals overlap to form delocalized pi-orbitals that extend throughout the sheet.
draw hectogon (6 dots each) shape with 5 connected. Create 3 rows of that. flat sheets. Looks like honeycomb.

Graphite: Leftover p orbitals form a band of pi molecular orbitals that extend throughout the entire structure so electrons are free to move. Delocalized electrons allow graphite to conduct electricity. Sheets of graphite are 2D and are held together by LDFs (relatively weak). LDFs are easy to break so sheets can slip and bend. Graphite is soft because of the slipping of the sheets.

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20
Q

Draw a molecular level picture of Diamond and use it to describe the bonding present

explain how the model of bonding you discussed explains the properties you listed for each substance

A

3D network of sp3 hybridized C atoms with localized covalent bonds.

High MP because covalent bonds are strong. Electrons are localized in the bonds between the atoms and cannot move freely. Therefore, diamond cannot conduct electricity. Diamond is hard because the atoms are covalently bonded in 3D structure, which is hard to break apart.

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21
Q

define a covalent bond

A

two atoms interact by ve- of one atom being attracted to the nucleus of another, however each nucleus is attracting the ve- in a tug of war of attractive forces.

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22
Q

define bond length

A

most stable distance between the atoms - lowest potential energy

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23
Q

when bonds form….

A

energy is RELEASED into the surroundings!

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24
Q

What is the molecular orbital (MO) theory?

A

electrons are waves and can combine constructively and destructively. In MO Theory, n atomic orbitals combine and give n molecular orbitals.

There’s a drawing of seven valence electrons for each of 2CL atoms to make a CL2 mole it is drawn into a molecular or remember to start at the bottom and work your way to the top!

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25
How do "bonding" molecular orbitals form? Do they have lower or higher energy then atomic orbitals? How do electrons behave in the species?
- when atomic orbitals combine constructively. - Lower -Makes it more stable
26
How does "antibonding" molecular orbitals form? Do they have higher or lower energy than atomic orbitals? How do the electrons affect the species?
when atomic orbitals combine destructively They have higher energy than atomic orbitals - electrons make species LESS stable
27
Draw a picture at the molecular level showing London dispersion forces between. Use the text box to explain what is causing the attraction between the atoms.
The picture shows originally a circle with the dots inside of it and a oval moving towards it with a dot dot in it. Then two ovals that are an induced staple that show the Delta charges plus and minus where the positive and negative side are next to each other. The textbox says: an atom that’s electron shift from one side to another becomes an instantaneous dipole. When this dipole gets close to other normal atom, the electron will then become attracted to the nucleus of the induced dipole. This then causes the atom to become attracted to each other as the opposite forces of both atoms are attracting each other.
28
What makes the oxygen molecule (O2) "paramagnetic"?
because of the way the electrons are arranged (it has unpaired electrons) - this is an example of how molecular level interactions determine the macroscopic properties of the substances.
29
Properties of Metals.
shiny, conduct electricity, malleable and ductile, can be grey colored, or colorless!
30
What makes a metal shine?
absorption of a photon will promote an electron to a higher energy level - if it immediately falls back down, emitting a photon, the metal shines!
31
what permits electrons to freely drift between bands?
over lap!
32
Explain in your own words why the ionization energy increases as you go from left or right across the periodic table.
As you go from left or right, the ionization energies increases. This is because the effective nuclear charge increases across a row, pulling the electrons closer to the nucleus and causing there to be stronger attractions that require more energy to be broken ionization energy increases (in general) as you go from left to right across the periodic table because more energy is required to pull the outermost electrons from those elements electron cloud because the smaller radius causes them to be closer to the nucleus. Being closer to the nucleus makes the attractive force even stronger which means more ionization energy will be required to break it.
33
The properties of the product of a chemical reaction are ....
emergent! we cannot predict the properties of a product from the properties of the reactants! they're unrelated!
34
Substances depend on the bonding/interactions that exist ____ the substance
within!
35
explain the similarities and differences in LDF and covalent bonds
London dispersion forces and covalent bonds have similar causes (electrostatic attraction of the electron of one atom to the nucleus of another). The differences are in: -The magnitude of the attraction -How the electrons are arranged in the new species - Covalent bonds are stronger, it’s hard to predict bond strength and its present only when atomic orbitals bond constructively and are present within molecules or networks.
36
Explain in your own words why hydrogen (H) atoms can from a covalent bond but helium (He) atoms cannot:
Hydrogen atoms can form a covalent bond while helium atoms can because the two electrons in the two hydrogen are in the bonding sigma with nothing in the anti-bonding sigma. However, with helium has four electrons between the two of them, which means as two and both the bonding Sigma and two in the bonding. This causes the bonding sigma to be canceled out and no covalent bonds to form.
37
What type of bond exist in the chlorine molecule?
single
38
difference between melting and boiling point?
When a substance melts some of the interactions between the particles must be overcome so that they can move relative to each other (what causes liquid to flow). When a liquid boils, all of the interactions between the particles must be overcome The magnitude of the melting point and boiling point provides an estimate of the strength of attractive forces between particles.
39
what is the difference between discrete and continuous materials? Give examples of each!
Discrete - exist as separate atoms or molecules Ex: He, Ne, Ar (any noble gas), H2O, H2,O2 Continuous - exist as extended networks of atoms connected to each other Ex: metals, diamond, graphite
40
What is the carbon bond like in diamond?
each carbon atom forms 4 bonds to 4 identical carbon atoms The bonds arrange themselves towards the corner of a 4 sided figure (a tetrehedron) with tetrahedral geometry C-C-C bond is 109 degrees.
41
What is the Valence Bond Theory?
To explain the bonding and carbon, we need to use the valence bond theory In this theory we assume that atomic orbitals overlap to form bonds Bonds are often shown as sticks “-“
42
How can carbon form four identical bonds (in diamond)?
(4 hybrid orbitals) ! They are as far apart as possible and the Valence Bond theory that for each bond carbon needs a singlily occupied orbital pointing in the direction of the bond. Sp3 hybridization s + p + p + p = sp3
43
what is a "sigma" bond?
Hybridized atomic orbitals (sp^3) overlap to give strong directed bonds causing a tetrehedral shape and “sigma bonds” Sigma bonds form when atomic orbitals overlap end to end. (cursive o) If diamond were to melt these bonds would have to be broken to allow them to move relative to one another which would require a lot of energy energy, therefore diamond decomposes instead of melting!
44
differences in valence bond theory and molecular orbital theory
Valence Theory - atomic orbitals overlap, the greater the overlap the stronger the bond, each bond is made up of two electrons, electrons are localized within the bond Molecular Orbitals- atomic orbitals combine to form an equal number of molecular orbitals each orbital can contain up to two electrons, electrons in bonding orbital stabilize the system , electrons in antibonding orbitals make it less stable, electrons are delocalized.
45
explain diamonds properties
Hard/High melting/boiling points - 3-D network of strong bonds! Does not conduct electricity - electrons are localized in bonds between atoms - not free to roam. There is a large band gap between the bonding and antibonding orbitals. Translucent - light passes through or is reflected. To absorb light an electron must be promoted to a higher energy level. There is a large “band gap” between the bonding and anti-bonding orbitals.
46
compare diamond and graphite!
Diamond has a high melting point. It’s hard, brittle, translucent and does not conduct electricity Graphite has a high melting point. It’s soft slippery and conducts electricity. However, both are pure carbon!
47
explain how the bonding in graphite is different than carbon
One S and 2 p orbitals hybridize to give us 3 sp^2 orbitals (with one P orbital left over). s + p + p = sp^2 The geometry is called trigonal planar. the CCC bonding is 120 degrees. When the SP2 hybrid orbitals combine, they form sigma bond molecular orbitals. The leftover P orbital(one on each carbon) combines side to side to form a large number (pi) of molecular orbitals!
48
explain graphite's properties
Conducts electricity - because electrons can move freely over the entire sheet within. It’s the localized pi molecular orbitals Shiny - because it can absorb a photons of many wavelengths (like metals) Slippery- sheets can slide over each other (only held together by LDF)
49
Draw and label an atomic/molecular level picture of the bonding in Bromine (Br2) that can be used to explain bromine's properties (mp 266k, brown liquid at room temp, low bp 332K...) Use your drawing to explain how the bonding and/or interactions in bromine lead to the observed melting point (266k).
picture shows a 3 x 3 row of bromine molecules (ovals with two dots to represent nucleus) in even, equal rows. Then it shows an arrow that says "melt" and then a disorganized set of (Br2) molecules in the shape of ovals still. explanation: Bromine is interacting with LDFs. Bromine has a very low melting point, only a little bit of energy is required in order for Bromine to melt. This means that there are weak interactions occuring. These weak interactions between molecules must be LDFs.
50
what are the steps to figuring out "isomers"?
1. Start w/ the longest carbon chain to draw the first isomer 2. Shorten the carbon chain by 1 carbon and put the remaining carbon in the middle of the longest carbon chain 3. Repeat until you cannot shorten the chain anymore!
51
How many isomers does C2H7N have?
2!
52
What is linear geometry?
bond angle - 180 degrees # of electron centers - 2 hybridization: sp
53
what is trigonal planar geometry?
of e centers - 3 bond angle - 120 hybridization - sp^2
54
what is tetrahedral geometry?
of electron centers - 4 109 degree bond angle sp^3 hybridization
55
how to calculate formal charge:
valence electron of free atom - # of bonds to atom in structure - # of non-bonded electrons on Atom in structure ex: (NH4) N FC = 5-4-0 =1 H FC = 1- 1-0 = 0
56
If you have multiple bonds...
they only count as one electron center in VSEPR
57
What is the overall difference between pi and sigma bonds?
pi bonds DO NOT rotate around the bond (it would break!) while sigma bond allows for rotation
58
what do CH4, H20, and NH3 have in common? Do you think they all have the same molecular shape?
4 centers of electrons tetrahedral geometry sp^3 hybridization - CH4, NH3, and H2O do not have the same shape, even though the electron center geometry is the same because there are a different number of bonds (sticks) AND the number of lone pairs is not counted in the molecular shape.
59
are electron center geometry and molecular shape the same?
NO! Geometry considers the positions of all the electrons while shape considers the position of the atoms (the bonding electron centers).
60
what is the shape of NH3?
trigonal pyramidal
61
shape of NH4+?
trigonal pyramidal
62
what is electronegativity?
the ability of an element to attract electrons to itself in a bond. An atom that has a high effective nuclear charge attracts its own valence electrons strongly and it attracts electrons from other atoms in bonds. electronegativity increases across (-->) the periodic table and decreases down the periodic table.
63
what does electronegativity depend on?
effective nuclear charge and size of orbitals
64
what is the molecular shape for CO2?
linear (2 centers due to double bond)
65
what is a "polar bond"?
when two atoms of different electronegativities bond, the electrons are not shared equally. This results in a bond dipole and the bond is polar. ex:/ HF you draw a cross arrow thing pointing towards the stronger bond.
66
What are the 5 steps to determine if a molecule is polar?
1. draw lewis structure 2. count the electron centers around the central atom to determine the electron center geometry 3. Determine the molecular shape 4. Determine bond polarities 5. Add up the bond polarities (take into account the direction - they are vector qauntities). If there is an overall diple, the molecule is polar. If not, the molecule is nonpolar.
67
What are the three types of intermolecular forces?
weakest to strongest: LDFs, dipole dipole, H-Bonds
68
What would you draw to show the IMF in liquid methane (CH4)
you would draw the lewis structure and then circle it and give it the delta plus and minus charges. Then show an arrow between the two draw and label LDF.
69
What IMFs are present in liquid F2?
LDF
70
what IMFs are present in liquid HBr?
LDF, dipole-dipole
71
How to draw dipole dipole interactions
draw two ovals with lewis structure inside and then show the cross dipole line thing pointed toward stronger molecule.
72
What IMF are present in liquid HF?
LDF, dipole-dipole, H-Bonding interactions
73
How to draw H-Bonding:
dashed line connecting H to a lone pair of electrons of another! cannot be connected to C!
74
Methanol (CH3OH) and Ethane (CH3CH3) have similar masses, what would you predict to have the higher boiling point? Why?
Methanol (CH3OH) Methanol has LDFs, dipole-dipole, and hydrogen bonds while Ethane only has LDFs. Hydrogen bonding is a much stronger interaction than LDFs and would take more energy to overcome the interactions. Therefore, methanol has a higher boiling point.
75
Which would have the higher boiling point? Methanol (CH3OH) or Ethanol (CH3CH2OH)?
Ethanol (CH3CH2OH)! Both compounds have LDFs, dipole-dipole interactions, and H-Bonds but Ethanol is larger and therefore will have more LDFs. Ethanol and Methanol both contain H-bonds and dipole-dipole interactions, but Ethanol has more electrons, so its LDF's are stronger.
76
Will a liquid sample of CH3NH2 display H-bonding interactions?
Yes! H-N
77
What type of IMFs are present in a liquid sample of CH2O?
LDFs and dipole-dipole interactions
78
How can you predict properties from a molecular formula?
Identification of individual bond polarities Identification of molecular polarity Deduce type and relative strength of intermolecular forces
79
What is the strongest IMF in He, H2, CH4, H2O, and NH3
He - LDF H2- LDF CH4- LDF H2O - Hydrogen Bonding NH3 - hydrogen bonding
80
How to identify if molecules are the same or isomers?
to determine if they are isomers, compare the molecular formula explain how the atoms are connected in the molecule consider bond types and geometry ex: image shows 2 as the central bond with OH, 1, 2, 3 being other "sticks" or bonds. There is another one but the shape is flipped, but 2 is still the central atom. they're the same! explanation: The –OH group is on the 2nd carbon in a 4-carbon chain
81
what is the difference in cis and trans isomer?
Isomers Double bonds do not have free rotation because the pi bond would break. Cis = same side Trans = opposite sides
82
Use Valence Bond Theory to explain why rotation is possible around single bonds, but not double bonds.
Single (sigma) bonds are formed when atomic or hybrid orbitals overlap directly between the two nuclei (end-to-end overlap of orbitals). When the bond rotates, the overlap stays the same, so the bond is not broken. Double bonds contain one sigma and one pi bond. Pi bonds are formed from the side-to-side overlap of parallel p-orbitals and result in the overlap occurring above and below the plane of the nuclei. If a pi bond were to rotate the overlap of the p-orbitals would be broken and therefore the bond would break.
83
What molecular shapes are possible starting from tetrahedral geometry?
Tetrahedral, Trigonal Pyramidal, and Bent
84
What molecular shapes are possible starting from trigonal planar geometry?
trigonal planar and bent
85
What is the difference between a trigonal planar and trigonal pyramidal molecular shape?
The trigonal planar shape is flat and the bond angle between all electron centers is 120 degrees. The trigonal pyramidal shape is a pyramid and the bond angle between all electron centers is about 109 degrees. The 4th electron center on the central atom in the trigonal pyramidal shape repels the other three electron centers and pushes them out of the plane creating the pyramidal shape.
86
Explain different structures of ionic compounds and use the structure to explain the properties of ionic compounds (for example melting point, boiling point, hardness, and ability to conduct electricity).
Ions form repeating 3D lattice structures Properties: strong ionic bond, shifting layers repel like charges, and conducts electricity only when molten or dissolved (ions move freely).
87
Explain why metals tend to form positive ions and non-metals tend to form negative ions.
metals have fewer valence electrons and when the lose them they become stable and therefore form positive ions whereas nonmetals have nearly full valence shells and therefore they gain electrons to become more stable and form negative ions
88
explain how to use Lewis structures and VSEPR to deduce electron pair geometry and molecular shape of molecules.
Lewis structures show bonds and lone pairs and VSEPR shows how electrons repel causes different shapes to form (linear, bent, tetrahedral) Use both to get geometry and molecular shape
89
Predict the polarity of bonds using atom electronegativity. Predict the polarity of molecules using bond polarity and molecular shape.
Bond polarity - use electronegativity difference (bigger the difference - polar bond) molecule polarity - look at bond polarities and shape.
90
Explain differences in melting and boiling points in terms of forces and energy.
Higher MP/BP - stronger forces more energy is needed to break stronger interactions.
91
Explain why metals tend to form positive ions and non-metals tend to form negative ions.
metals form positive ions because they have few valence electrons and low ionization energies, so it's easier for them to lose electrons to reach a stable (noble gas) configuration. In contrast, nonmetals form negative ions because they have high electronegativity and more valence electrons—so they tend to gain electrons to complete their outer shell and become more stable.
92
. Explain why a third body (another atom or molecule) is (almost always) needed to form a stable bond.
when a bond is formed, it releases energy and it needs somewhere to go the 3rd body absorbs the excess energy and stabilizes the system.
93
What determines whether a bond or interaction is stable? Explain how temperature influences the stability of bonds and interactions.
stable bonds have low potential energy, and strong attractions higher temperatures add energy and can break weak bonds.
94
Explain on the atomic/molecular level why metals are malleable, ductile, and conduct electricity.
metals atoms are a sea of electrons electrons move freely therefore conducting electricity the layers can slide and move causing it to be malleable and ductile.
95
explain why relative melting points and boiling points for substances that exist as molecules (like H2) differ from those that exist as continuous extended networks (like diamond or metals).
molecules have weak forces and therefore low MP/BP whereas extended networks have a strong continuous bond and therefore require a higher MP/BP.
96
How does ionization energy support the idea that electron energies are quantized?
If electrons could exist with any energy level (that is if their energies were not quantized), any amount of energy might remove an electron. The fact that ionization energies are constant and replicable for any given electron means that this amount of energy must be added to the atom to eject the electron. (This is a similar argument to the idea that emission and absorption spectra are evidence for quantized energy levels – except here the electrons are ejected from the atom rather than moving between energy levels.)
97
Out of Ar and CL2 which will have the higher boiling point?
CL2 because when this substance boils only LDFs are overcome and the strength of the LDFs depends on the size of the electron cloud.
98
Draw the Lewis structure for sulfur dichloride, SCl2. How many lone pair of electrons are on the sulfur atom?
2 lone pair
99
A Lewis structure for nitrous oxide (laughing gas) is shown to the right. What is the formal charge on the nitrogen atom in the middle of the molecule?
+1
100
when Oxygen is the central bond connected to H and C, what is the center geometry?
tetrehedral
101
15. The Lewis structure for ethene (C2H4) is shown below. What is the hybridization and the shape around the carbon labelled with the arrow?
sp^2 and trigonal planar (Carbon double bonded to c on the right and bonded to two bent H's)
102
Which element is more electronegative N or P?
N - because the bonding electrons are closer to the nucleus
103

CEM 141 (3 decks)

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