atomic structure review

Question 1

J Dalton

Answer 1

• All matter is made up of tiny particles called atoms. An atom cannot be created, destroyed, or divided into smaller particles (not true because the atom is made up of p,n and o).
• The atoms of one element cannot be converted into atoms of another element.
• All atoms of one element have the same properties, such as size and mass. These properties are different from the properties of other elements. (not true because we have isotopes which have diff masses but same element)
• Atoms of different elements combine in specific proportions to form compounds.
• Billiard ball model (dense and thick atom)

Question 2

JJ Thompson

Answer 2

• Using cathode ray tubes it was demonstrated that the atom could be broken down into smaller particles.
• Negatively charged particles(electrons) could be ejected from atoms, leaving the atoms positively charged.
• raisin bun model (neg charge is rasins in a bun which is sphere of pos charge)

Question 3

Ernest Rutherford

Answer 3

• Directed highly energetic, positively charged alpha particles at a very thin gold foil.
• The pattern produced by the alpha particles proved that there was a dense positive core, this core contained almost the entire mass of the atom.
• It was inferred that the electrons exist in orbit around the small nucleus of the atom.
• since it glowed everywhere it indicated that the pos charge is small in the middle (pos alpha particles deflected slightly when near nucleus, but deflected at large angle when hitting nucleus directly which was not the expected result of the test)
• behive model (the nucleus is the hive and e- are bees buzzing around)

Question 4

N Bohr

Answer 4

• Analyzed the pattern and colours of light emitted from heated hydrogen atoms.
• Determined that the light being observed was being caused by electrons which were transitioning from a “higher orbit” to a “lower orbit”, thereby releasing energy.
• Developed a mathematical expression to calculate the radius of specific allowable energy levels where electrons can exist in orbit around the nucleus of the atom.
• excitation it takes in energy and deexcitation is gives off light or energy –> colours are diff because e- on specific levels will release diff colours based on where it moves down from
• planetary model (atom looks like the planetary system in which the nucleus is the sun and the e- are like planets orbiting)

Subatomic structure: the average atom is 10^-10 m in diameter but the atoms are made up of many smaller subatomic particles (p and n have relative mass of 1 while e- has 1/1840)

Question 5

Electromagnetic radiation

Answer 5

James Clerk Maxwell in 1856 proposed the existence of waves which were related to both electricity and magnetism and called them electromagnetic waves. (Electromagnetic wave has an electric and magnetic field component. Both have same wavelength and frequency.)

on the spectrum it goes from low freq and long wavelength to high freq and short wave length –> wavelength and freq occur together –> c which is speed of light = 3.00x10^8m/s –> c=wavlength * f –> f(sec-1) = c/wl –> wl= c/f

Some waves require a medium to travel through. for example water waves and sound waves.
in a vacuum no sound waves will be heard.

ELECTROMAGNETIC WAVES, requires no medium of transmission. Waves will travel in a vacuum

Question 6

max Planck

Answer 6

• Founder of the quantum theory
• Proposed that radiation is emitted in discreet packets called quanta (singular quantum)
• Energies associated with these quanta are proportional to the frequencies of the emitted radiation.
• He formulated the following equation: E (J) = hf
  Where E is the energy of radiation, h is Planck’s constant (6.626 x 10^-34 Js) and f is the frequency of the radiation.
• another equation could be E= hc/wavelength –> so e is indirectly proportional to wl (as one increases the other decreases)

Radiation emitted by heating solids to various temperatures. The energy emitted by atoms or molecules is always in whole number multiples of energy of certain well-defined quantities. The word Quanta or Quantum refers to the smallest quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation. E=hν, where h is Planck’s constant = 6.63 x 10-34 Js

Question 7

Einstein

Answer 7

• Theory of relativity
• Theory of the photoelectric effect – Noble prize in 1921
• Einstein viewed electromagnetic radiation as beams of photons
• each photon is a little packet of energy with the value E = h f.
• Frequency is equal to the speed of light divided by the wavelength of the radiation in question and thus this equation can be rearranged to read:
  E = h c
  λ
  Note: The theories of Planck and Einstein were the first steps in the development of the quantum theory.
• he basically was saying that waves are particle like

Proposed theory to explain the photoelectric effect. Einstein’s theory of light states that a beam of light behaves as a stream of particles rather than wavelike. He called the stream of particles PHOTONS. Later it was proposed that light behaves as either as waves or particles. This property is also a characteristic of all matter.

Question 8

DeBroglie

Answer 8

• French physicist
• Received the Nobel Prize in physics in 1929
• Louis de Broglie proposed the wave-particle duality of light. This simply put, is the phenomenon that waves behave both like a particle and as a wave (everything has wave like properties but smaller stuff has more)

De Broglie proposed that both light and matter obey the equation:

λ = h / m v
Where: λ is the de Broglie wavelength, m is the mass, v is velocity, h is Planck’s constant

If waves can behave like a stream of particles (photons) then so can an electron possess wave like properties. A particle in motion can be treated as a wave and vise versa.

Question 9

back to bohr

Answer 9

• Danish physicist
• First quantum model of the atom
• Formulated a description of the hydrogen atom that explained its light spectra
• Using Einstein’s work and Planck’s work – Bohr related the colour of hydrogen’s spectra to wavelength
• Wavelength to energy
• Bohr postulated that the colours of the Hydrogen spectra were discreet bands of colour.
• Bands generated by de-excitation of electrons from higher energy levels to lower or ground state, while emitting a photon of light
• Bohr showed that electrons are quantized (have specific values).
• He demonstrated that the only possible energies of electrons in his orbital/energy levels would be given by the equation:

    En =	- 1311 KJ/mol
            n^2		

The En values corresponds to the energy states of electrons in a hydrogen atom.

Hydrogen spectrum of light:
- When an electron makes a transition from a (excited state) higher energy state to a lower (ground state) energy state, it is accompanied by the emission of a photon having a wavelength that corresponds with the electromagnetic spectrum.
- When electrons are excited in a gas sample or element, radiation is emitted at discreet wavelengths.
- This is called a line spectrum.
- If we are examining a gas consisting of individual atoms then we call the spectra atomic emission spectra.

Question 10

J Balmer

Answer 10

• an amateur Swiss scientist
• developed a mathematical relationship that could explain and predict the visible spectrum of hydrogen.
• Johannes Rydberg – a Swedish physicist - using the work of Balmer developed an equation that related all the wavelengths of lines produced in the electromagnetic spectrum of Hydrogen. The equation is as follows:

1/wl = Rh (1/n1^2 - 1/n2^2)

RH (Rydberg constant): 1.097 x 107
n1 or nf: is the series type (final energy level)
n2 or ni: is the energy level the electron is transitioning from (initial energy level)

The hydrogen electron has a variety of transitions from different energy levels. The line spectra that we see is only in visible range. However, there are transitions that occur in ultra violet and infrared region of the spectrum.
- electron absorbing photon goes up energy level, electron emitting photon goes down energy level
- going the n = 1 is uv light (layman series, going to n=2 is visible spectra (balmer series), going to n=3,4 is infared light (paschen series, brackett series)

• the closer to the nucleus, the more negative you are, there are infinity n

Question 11

Spectroscopy

Answer 11

• A continuous spectrum is produced when white light is passed through a prism. The result is line a rainbow, where ROYGBIV can be seen.

Absorption Spectra:
In order to excite an electron a
certain amount of energy is required (at a particular wavelength) corresponding to the differences between the energy levels. For example, an object that absorbs blue, green and yellow light will appear red when viewed under white light.

Emission Line Spectra:
This is the opposite of absorption line spectra. The energy released when an electron falls back down to the ground state (at a particular wavelength) corresponding to the difference between
the energy levels.

Uses: to identify elements present in a gas or liquid, for example elements in stars and other gaseous objects which cannot be measured directly.

Question 12

Bohr and imrpovment of atomic model

Answer 12

• The number of electrons at each principle energy level became known
• # electrons = 2n^2 where n is the principle energy level
• evidence/reasoning ascertained from the intensity of the spectral lines of hydrogen.
• Electrons must be located in levels of specific and fixed energies (orbits).

    En =	- 1311 KJ/mol
            n^2	   Evidence/Reasoning – given that hydrogen gave off four very distinct bands of colour, where the photons must have discreet amounts of energy, since this energy was released when excited electrons dropped from higher energy to lower energy, the electrons must have specific energy levels at which they must exist and not in between.

limitations:
- The electron is a particle whose position and motion can be specified at a given time.
- An electron moves in an orbit having a fixed radius.
- Bohr’s experimental evidence only agrees with an one-electron atom (Hydrogen)

… in addition to this, electrons behave as particles in some experiments while in other experiments they behave as WAVES

(Atomic structure was analogues to planetary system. Studied line spectra of hydrogen in gas phase. Line spectra due to the excitation of electrons from lower energy level to a higher one. When the electron falls to a lower energy level or ground state it emits a specific amount of energy. The line spectra are not restricted to the visible range but also extend into the UV and IR range. The energies associated with electron motion in permitted orbits are fixed in value. Therefore energies are QUANTIZED)

Question 13

E Shrodinger

Answer 13

• Developed a mathematical expression called the “Schrodinger Wave Equation”, which when solved, predict the CHANCE or probability of an electron showing up in a particular region of space (3-D) around the nucleus.
• this calculated region of space is called an Orbital. E
• Orbitals are where electrons are confined to specific regions, which are represented as electron clouds.

In order to describe the motion of an electron in an atom, both the electron’s wave and particle nature must be taken into account.

HEISENBERG UNCERTAINTY PRINCIPLE states that due to the dual nature of the electron it is impossible to determine both its position and energy at the same time.

Question 14

Probability model of the atom

Answer 14

1. In this model, the probable location of an electron moving about the nucleus can be identified, but not its exact location. The general location or ‘region of space’ within which the electron can be found is called an orbital. No longer will the electron be thought of as travelling in a defined path called an orbit.
2. The shape of the orbital or region of space will be defined by an array of dots where each dot represents a possible location of the electron. The more intense the dots the greater the probability of finding an electron at that position. The diagram above results from a computer program that has determined the possible positions of a hydrogen atom’s single electron.
3. The maximum capacity of any orbital is TWO electrons, each having opposite spin. Since electrons are negatively charged, their rotation produces a magnetic field. Therefore, opposite rotation produces unlike/opp magnetic field which in turn produces the attraction to allow the two electrons to exist in the same orbital.

Question 15

Types of orbitals

Answer 15

• S orbital: this is a spherical shaped region of space –> only one orietntion in space –> max capacity of 2 e
• p orbital: dumbell shaped region of space, 3 orientations, so full set has a max capacity of 6 e
• d orbital: complex shapes-4 lobes mostly, 5 orientations so full set has capacity of 10 e
• f orbital: complex shapes, 7 orientations so full set has capacity of 14 e

When scientists studied multi-electron systems (elements other than hydrogen), they discovered that there were more levels (more frequencies) visible. The basis for these sublevels/orbitals is due to studies of spectra of different elements that indicated that each shell or energy level consists of one or more orbitals/sublevels grouped closely together.

n = 1, s = 2e
n = 2, s =2e and p= 6 e
n = 3, s = 2e, p= 6e and d = 10e
n = 4, s=2e, p=6e, d= 10e and f = 14e

orbital energy/tables are drawn in increasing energy going up, have to spread out arrows before doubling

electron configurations: Electron arrangements can be made more specific than the Bohr-Rutherford model has to date. This model is based on the quantum mechanical model. It specifies the level the electrons are on and is referred to as electron configurations. As with the Bohr model though, electrons always fill the lowest energy state before proceeding to a higher energy state, regardless of the actual being occupied. Follow the pattern on the right for the order of filling up the sublevels.

ex. 1s^2 2s^2 2p^6 for neon where first number is energy level then orbital type and then number of e there

AUFBAU PRINCIPLE – electrons go into the lowest energy level and fill them in order of increasing energy.

PAULI EXCLUSION PRINCIPLE – an orbital can hold a maximum of two electrons. But an orbital can be empty or have one electron. (0,1,2)

HUND’S RULE – electrons in the same sublevel or orbitals will not pair up until all the orbitals of the sublevel are at least half-filled.

ELECTRONIC CONFIGURATION – distribution of electrons among the various orbitals.

Question 16

extra info

Answer 16

Davisson, Germer, G.P Thomson:
Experimentally proved that electrons possess wave-like properties.

Werner Heisenberg:
Heisenberg uncertainty principle states that it is impossible to know both the momentum (speed x mass) and position of a particle with certainty.

Erwin Schrodinger (1926):
Formulated an equation that would describe the behaviour and energies of submicroscopic particles in general. Used this equation to find the probability of locating an electron in a given volume. This led to quantum mechanics or wave mechanics model.

Quantum Numbers:
Describe the distribution of electron in an atom

Principal quantum number:
Relates average distance of electron from nucleus to a particular orbital. The bigger the number the larger the orbital. Energy level (n)

Angular quantum number:
Sub shells. Orbital shapes (l). s, p, d, f, g, h…

Magnetic quantum number:
Orientation of orbital in space (ml). The number of orientations can be calculated using formula 2l+1

Electron spin quantum number:
Orientation of electrons within a subshell (ms). Two electrons is the same orbital will have opposite spins due to their opposite magnetic fields.

Electronic configuration:
Number of electrons distributed among various atomic orbitals.

Pauli Exclusion principle:
No two electrons can have the same four quantum numbers. Therefore electron spins must be opposite.

Hund’s rule:
Most stable arrangement of electron in subshell is one with greatest number of parallel spins.

Aufbau principle:
Aufbau means building up in German. The process of building up the ground state structure for each atom, in order of atomic number.

Ionization energy:
The energy needed to completely remove an electron from a ground state gaseous atom.

Question 17

ec or energy level diagrams or pt

Answer 17

Electron configuration is when you write out 1s^2 etc but it also is the chart/model that specifies the level that the electrons are on –> you write out 1s to 7s, 2p to 7p, 3d to 7d and 4f to 7f and then you draw diagonal lines to see the order of the erngy levels/orbitals

you can also look at the periodic table to find out the levels –> the first two families are the s orbital, groups 13-18 are p orbitals, the transition metals are d and the 2 series at the bottom are both f

energy level diagrams show the spin of the electrons on every energy level (shows that 1s is however many half arrows etc).

Question 18

Electron configuration

Answer 18

Electronic configurations for a particular atom or ion can be expressed in three ways:

Orbital box diagrams (or orbital diagrams) – a physical box (or circle) represents the orbital and electrons with spins pointing up or down are placed within the orbital.

Electronic configuration – uses the energy level, orbital letter and the number of electrons within the orbital is expressed as a superscript. (ex. 1s^2 etc.)

Condensed or Shorthand electronic configuration – the noble gas in square brackets used to represent the inner electrons followed by the electronic configuration of electrons in the outer shells of the atom. Remember in chemical reactions inner electrons are not involved; it is the outer valence electrons that take part.

But there are expectations to this rule. Chromium does not work as instead of 4s^2 3d^4 it becomes 4s^1 and 4d^5 as it gains half filled stability. Copper also goes from 4s^2 4d^9 to 4s^1 4d^10 as it also gets half stability in s and full stability in d. this rule is that filled and half filled subshell have extra stability that sometimes affects the electron configuration.

S block is group 1 and 2 (exception of helium). D block is transition metals. P block is group 13-18. F block is the series at the bottom.

Deviation from the trends: Ionization Energy (kJ/mol)

X (g) 🡪 X+(g) + e-

I.E. is the energy required to remove an electron an isolated, gaseous atom. In the periodic table there is a general trend in ionization energy. IE increases from left to right and decreased from top to bottom. However, in the ionization energy trend there are some values that have variations. Ex. Be has a higher IE compared to B this is because Be has full stability, the 2s is lower energy to more energy required to remove the e as its closer to the nucleus whereas in B the 2p level is higher in energy and further from the nucleus causing it to be easier to remove an electron.

Question 19

General terms of the quantam mechncial model

Answer 19

ground state electron configuration as the normal configurations that can be made from the periodic table. Excited electron configuration is when an e from the last orbital in the last energy level is moved up energy levels example from 3s^2 it becomes 3s^1 5s^1

Isoelectronic is when species have the same electron configuration meaning they have the same # of e- in the same places (configuration matches) ex. ne and fe-1

Energy level is the energy associated with a particular shell in which electrons reside. Orbital is a region of space where there is a high possibility of finding an electron. Orbut is the elliptical path where electrons were thought to exist.

x,y is where an orbital lies

Question 20

Quantam theory

Answer 20

Werner Heisenberg

German physicist
1901 – 1976
Leader in the development of quantum theory in the 1920’s

Heisenberg carried out a careful analysis, which showed that it is not possible to determine as electron’s momentum (mass x volume) and its position/location simultaneously. IF we know one we cannot know the other. This is known as the Heisenberg’s uncertainty principle: it is impossible to determine the location and momentum of an electron simultaneously. But electrons can be described as being located in orbitals, which are 3d spaces where there is a high probability of an electron being found.

The size, shape and orientation of these orbitals are determined by solving Schrodinger’s wave equation.

The exact solution of the equation yields the four quantum numbers. These numbers are the electrons ‘address’. No two electrons in the atom have the exact same set of quantum number.

Question 21

Quantam #’s

Answer 21

Quantum numbers are needed to describe distribution of electron. There are three quantum numbers needed to this: Principal quantum number, Orbital-shape quantum number or angular momentum quantum number, magnetic quantum number, spin quantum number.

Principle Quantum number (n): (energy level)
Indicated the size of the orbitals since it relates average distance of electron from nucleus in particular orbital. The bigger the n number the further away from nucleus an electron is.
The number of electrons in the energy level can be determined 2n2

Orbital-Shape Quantum Number aka Azimuthal number aka orbital angular momentum quantum number(l): (orbital type):
Indicated shape/type of orbital
l has possible integral values from 0 to (n –1):
l=0: s orbital
l=1: p orbital
l=2: d orbital
l=3: f orbital

Magnetic Quantum Number (ml): (diff position of the orientations)
Indicates the direction/orientation of orbital in space.
Indicates the number of orbitals in a subshell with a particular l value.
Total number of orientations can be calculated using the formula (2l+1).
Orientations can also be used by following this sequence: -l, (-l+1), …0, … (+l –1), +l or more simply integers from –l to +l. ex. l = 0 then ml = 0 but when l = 1 (p orbital) ml= -1,0,1 as it goes from -l to positive l

Spin Quantum Number (ms)

Evidence for spin was based on the fact that lines of spectra split in the presence of an external magnetic field.
Electrons act like tiny magnets, as the electrons spin on their own axis it generates a magnetic field of its own.
In 1924, Stern and Gerlech proved the electron spin nature.
Values are +1/2 or –1/2 (pointing up pos, pointing down neg)
Two electrons sharing a single orbital must have different spins (Pauli exclusion principle).

ex. the quantum numbers for the highest energy electrons in aluminum (3p1) would be:
n = 3
l = 1
ml = -1 (could be other number like 0 or 1 as well since no doubling)
ms = +1/2

Question 22

Valence bond theory

Answer 22

Based on wave mechanics it is possible to explain the bonding behavior in covalent molecules. Therefore, the use of atomic orbitals is essential to the bonding process and the electron location. There are two theories of covalent bonding: valence bond theory and molecular orbital theory

Valence Bond Theory:
Overlapping of two atomic orbitals such that the orbitals share the same region in space and only two electrons can be shared between the two orbitals.

Essentially, when two atoms like hydrogen come together there is a precise distance between the two orbitals that ensures maximum overlap of the two orbitals. The need for maximum overlap is responsible for the different shapes of molecules found in nature. The distance between the two nuclei in a bond is referred to as bond length. Remember the two shared electron of opposite spins spend MOST of their time between the two nuclei.

Overlap of orbitals can be between like or unlike orbitals.

When s and s bond then its one huge blob, s and p is one large block and one smaller one, p and p create a weird bond like e shared between one side of x axis orbital and one between the y axis orbital

The valence e- determine how they will bond and with what type of orbital which is why its necessary to write out the configuration.

However the VBT does not work with all bonds as it can not account for the bond angles found in nature, only for small diatomic molecules it can as it does 90 degree bonds.

Hybridzation theory is used as it uses orbitals and overlaps them in a way that they meet the bond angles and shapes of molecules found in nature.

Question 23

Hybridization Theory

Answer 23

Hybridization theory – is based on the mixing of like orbitals (s and s or p and p) or unlike orbitals (s and p or p and d etc) to create “hybrids”. Get it! hybridization = hybrid orbitals.
Note: the hybrid orbitals are different from the orbitals in the free state of an atom.

Use this when bonding two or more things

produces 1 type of shape which is a dumbell with one fat end or shaped like a balloon

Hyrbidization allows for the actual bond angles to form that can not be done by the VBT

You basically are making hybrid orbiatls that are all equivalent meaning that you mix two orbitals so they get to one level. This is done for the central atom. Only use valence electrons for this.

BOND TYPES: sigma bond which is the end of end overlap of atomic orbitals or hybrid orbiatls (single bonds where the elctrons on inter nuclear axis). Pi bond which is for double or triple bonds. Electrons density is above and below the bond axis so less strong because the bonding is far from the nucleus? and easier to break therefore

So how do I figure out the hybridization of the CENTRAL ATOM – anything bonded to two or more atoms
- Count the number of substituents attached to the central atom (count lone pairs are a substituent). This total number of substituents should equal the hybridization number.
- when considering the hybridization state of an atom, you must remember that only one s orbital can be involved in hybridization, similarly, a max of 3 for p and a max of 5 for d can take part in hybridization. (s1p3d5 : maximum allowed)

Steps to successful hybridization:
- Count the number of substituents attached (including lone pairs) of central atoms and determine the hybridization designation. For example, sp3, sp2, sp, sp3d1, sp3d2, sp3d3, sp3d4, sp3d5.
- Then write down the electronic configuration of the atoms involved.
- Use the valence electrons of the central atoms to form hybrid orbitals. The orbitals involved in the hybridization will be determined by step 1. For example, in sp hybridization, you will need one s orbital and only ONE p orbital. Since there are a total of 2 atomic orbitals going in then there should be the same number of hybrids being produced, therefore 2 sp hybrid orbitals will be produced.
- Draw the hybrid orbitals, any hybrid orbital with a single electron has the ability to bond with a substituent. If there are pairs of electrons left over, these will be the lone pairs on the molecule.

NOTE: there should NEVER be SINGLE electrons anywhere. If so, always go back and check your work.

(basically you draw the lewis structure and determine the hybridization designation based on the substituents and/or lone pairs, you then write the e- configuration, then draw the regular orbitals of the central atom and then combine them to form the hybridization and write how many hyrbid orbiatsl were formed and its designation, then draw the actual orbitals by starting with the central, draw the different hyrbid orbitals and bond with the substituents) (the amount of lines(based on designation) and e- that go in must come out)–> hybridization designation just tells you how many orbitals to take out ex. if sp2 then take out s and 2 p lines

you can break the octet rule according to the e- each orbital can hold and bohrs equation (2n^2) –> this can only be done by period 3 elements and beyond –> ex. PCl5 works –> you tally the amount of e-, draw the diagram, create the hybridization deisgnition (includes d in it), create the hybridization orbitals and then the diagram (if connecting to p draw the p orbital –> means that if the central atom is bonded to another atom that is not hybridized and has p orbital as highest energy make sure you draw p orbital not s)

with db and tb the approach is the same but sigma and pi bonds can form
steps of doing:
- draw the diagram
- create hybridization designation (based on # of things central atom(s) are connected to including lone pairs)
- write the e- configuration
- draw the hybridization orbitals of the central atoms, draw the orbital overlap diagram including any db or tb
- there may be some e- not included in the hybridization ex. some e- promoted are not included in the hybridization orbitals and are added on the side as x,y or z p orbitals –> these are then drawn for the double or triple bond so if only one side p orbital then a double bond but if 2 side p orbitals then its a triple bond (they carry only one e- and are bonded to another central atom that has a side p orbital with one e- to form a bond that contains 2e-) –> pi bonds have to be in same direction, either py or pz and have to be labelled

make sure that one e from the 2s orbital is promoted in the orbital drawings for hybridization

unhybridized doesn’t show lone pairs

(make sure to count lp for hybridization designation, e- promo happens form any level)

Question 24

VSEPR Theory 1

Answer 24

Valence Shell Electron Pair Repulsion

The rule is that like charges repel each other. Thus, within a molecule, electrons pairs will orient themselves such that there is maximum separation between them, which will lead to minimal repulsions. This desire to achieve minimum repulsions between electron pairs results in various shapes of molecules.

(getting as far as possible results in stability)

The VSEPR model was developed by Ronald J. Gillespie and R.S. Nyholm and is based on 3 main ideas:
- Chemical reactions and bonding involve only the electrons in the outermost or valence shell of an atom.
- Electrons in orbitals and in bonds are always in electrons pairs.
- Electrons repel one another because of like electrical charge.

Terminology: Sometimes the molecules are represented by AXyE, where Y is the # of peripheral atoms (things attached). A: central atom, X: bonded atoms, E: non-bonding electron pairs. Bonding pair (BP): 2 electrons involved in a bond (X). Lone pair (LP): 2 electrons NOT involved in a bond (E). Ex. AX3E1

Rules for using VSEPR theory:
- each pair of electrons in the valence shell occupies its own region of space called the domain of the electron pair.
- The bonding pairs (BP) and lone pairs (LP) move as far apart as possible to minimize electrostatic repulsions.
- A LP occupies slightly more volume than a BP
- Multiple bonds (double and triple) occupy more volume than a single bond.
- Polar bonds occupy less space at the central atom than non-polar bonds.

Note: VSEPR is only applicable to molecules or molecular ions. It does not apply to ionic compounds.

Types of shapes include linear (bonded to two, degree 180), trigonal planar (bonded to 3, 120 degrees), tetrahedral (bonded to 4, 109.5 degree), trigonal bipyramidial (bonded to 5 things, 90 and 120 degree) and octahedral (bonded to 6 things 90 degrees) –> axial is at top and bottom, while equatorial is in middle

Question 25

VSEPR Theory 2

Answer 25

Electron Pair Geometry (EPG): total number of electron pairs about the central atom (BP plus LP)

Molecular Shape OR molecular geometry: number of BP versus LP (looking at squishing bond effect of lone pairs)

• Lone electron pairs in the valence shell affect the shapes of molecules.
• The lone pair is more spread out – it is bulkier than a bonding pair of electrons and takes up more room than a covalent bond does. A lone pair also affects the bonding angles. It decreases the bond angles by roughly 2 degrees.
• Electron pair repulsion magnitudes decrease in the order: LP – LP > LP – BP > BP –BP
• VSEPR theory is applicable to molecules with multiple bonds. We count a double or triple bond as one group of electrons connecting the ligand X to the central atom A. Although, multiple bonds repel single bonds and thus multiple bonds act like lone pairs. Generally, we are going to treat a double or triple bond as a single electron pair and predict its shape using the rules applied for single bonded species.
• Lone pair electrons occupy the equatorial faces of trigonal bipyramidal molecule. Placing the lone pair at an equatorial vertex rather than an axial vertex minimizes repulsion due to lone pairs.
• Two lone pairs occupy opposite vertices in an octahedron. To minimize the lone pair repulsion in AX4E2 the lone pairs are placed at opposite vertices.

What you do is you draw the lewis structure based on the compound formula. Make sure you draw correct EPG shape based on the # of things attached to central atom including any lp (lp can be determined by how many valence e- exist and if they are all being used or not). Then write the EPG and the bonding and non bonding electron regions. Based on that write the molecular shape (changes if there any lp).

(based on total bonding regions (non and yes))
2 bonding regions –> linear, 180 degrees, linear ms
3 bonding regions –> trigonal planar, 120 degrees, 3bp then same but if 2 bp and 1lp then bent
4 bonding regions –> tetrahedral, 109.5 degrees –> 4 bp then same, 3 bp and 1 lp then trigonal pyramidal, if 2bp and 2lp then bent
5 bonding regions –> trigonal bipyramidal, 90 axial and 120 equatorial –> 5bp then same, 4bp and 1lp is seesaw, 3bp and 2lp is t shaped, 2 bp and 3 lp is linear
6 bonding regions –> octahedral, 90 degree –> 6 bp then same, 5 bp and 1lp is square pyramidal, 4 bp and 2lp is square planar , then t shape and linear again

domains for non-bonding e- pairs exert greater repulsive forces on adjacent e- domains and thus tend to compress the bond angles.

you can apply vsepr to complex molecules by creating the epg and ms for each central atom (things bonded to 2 or more things not rlly including lp) and then applying those shapes to draw it out

Question 26

Lewis Structures

Answer 26

• Sum valence e- from all atoms in the molecule (don’t forget about charges on ions!).
• Bond ligands to central atom (usually the atom that comes first in the formula; or the least electronegative atom) with a single bond.
• Complete octets around ligands.
• Place any leftover e- on the central atom (even if it results in more than an octet).
• If there is not enough e- for an octet around the central atom, try multiple bonds by moving a non-bonding pair of e- from a ligand to a position where it can be shared.

The first step to predicting the three-dimensional shape of a molecule is to draw its two-dimensional Lewis structure. Drawing a Lewis structure for a molecule lets you see
exactly how many electrons are involved in each bond, as well as how many, if any, lone pairs of electrons are present.

Step 1 Position the least electronegative atom in the centre of the molecule or polyatomic ion. Draw a skeleton structure for the molecule by placing the other atoms around the central atom. Draw a single bond between each pair of atoms. Always place a hydrogen atom or a fluorine atom at an end position in the structure.
Step 2 Determine the total number of valence electrons (V) in all the atoms in the molecule. For polyatomic ions, add or subtract electrons to account for the charge. For example, for a cation with a charge of 2+ subtract two electrons from the total number of valence electrons that are calculated for the ion.
Step 3 Determine the total number of electrons (T) needed for each atom to achieve a noble gas electron configuration. For all atoms besides hydrogen this corresponds to the octet rule. For hydrogen, a complete valence shell is two electrons, not eight.
Step 4 Subtract the number of valence electrons (V) from the number of electrons needed to satisfy the octet rule (T). This
represents the number of shared electrons (S) involved in bonding, S = T – V. Divide this number by 2 to give the number of bonds, bonds = S/2. Double bonds count as two bonds. Triple bonds count as three bonds.
Step 5 Subtract the number of shared electrons from the number of valence electrons to get the number of non-bonding electrons, NB = V – S. Add these electrons as lone pairs to the atoms to achieve a noble gas electron configuration for each atom.

if it says lewis dot diagrams create everything as dots including the share e- –> make them horizontal dots in pairs of 2)

Question 27

Formal Charges

Answer 27

Occasionally we can draw more than one Lewis structure for a molecule. So how do we decide which is the correct structure or stable structure?

Formal charge of an atom: is the charge that an atom in a molecule would have if all atoms had the same electronegativity.

Formal charge = [# of valence electrons of atom] - [# of electron assigned to the atom*]
or # of valence - dots - lines

• this can be determined by adding all the lone electrons and the half of all bonding pairs

Electron arrangement is best if:
The atoms bear the smallest formal charges. Any negative changes reside on the more electronegative atoms.

NOTE: Formal charges are NOT real charges or oxidation number.

• All unshared e- are assigned to the atom on which they are found.
• For bonding e-, half of the bonding e- are assigned to each atom in the bond.
• Subtract the number of e- assigned to the atom from the number if valence e- in the isolated atom.

Note 🡪 formal charges on neutral molecules sum to zero and in ions they sum to the charge on the ion.

Exceptions to the Octet Rule
🡪 3 main exceptions

• Molecules and polyatomic ions containing an odd number of e-.
  When the total number of valence e- is odd, complete pairing of e- is impossible and an octet around each atom is impossible. Example: nitrogen monoxide (N=O)
• Less than an octet of e-.
  Most often encountered with beryllium, boron, and aluminum. Example: boron trihydride (BH3). A boron atom only has 3 electrons in its outer level, and there is no possibility of it reaching a noble gas structure by simple sharing of electrons. Is this a problem? No. The boron atom has formed the maximum number of bonds that it can under the given circumstances, and this is a perfectly valid structure. Energy is released whenever a covalent bond is formed. Because energy is being lost from the system, it becomes more stable after every covalent bond is made. It follows, therefore, that an atom will tend to make as many covalent bonds as possible. In the case of boron in BH3, three bonds is the maximum possible because boron only has 3 electrons to share.
• Expanded Octet (hypervalent species): More than an octet of e-.
  Can only occur with elements of the third period and beyond. An expanded octet also depends on the size of the central atom, the bigger the size the higher the chances that the an expanded shell can be accommodated. Expanded shells often occur when the central atom is bonded to small and highly electronegative atoms like F, Cl, O and so on.** Example: SF6

Why do some elements exhibit expanded octets? elements of the third period and beyond have empty d orbitals present in their atomic structure. The empty d orbitals can accept more than an octet of e-.
Therefore, since sulfur only has e- placed in the 3s and 3p orbitals, it has space to place extra e- in its empty 3d orbitals.

Question 28

Co-ordinate (daltive) covalent bonds

Answer 28

co-ordinate covalent bond - a covalent bond in which one atom contributes both electrons to the shared pair of electrons

In a covalent bond, a pair of e- is shared between two atoms, each atom having contributed one electron to the bond. In some cases, one atom contributes both e- to the shared pair. This occurs when a filled atomic orbital overlaps with an empty atomic orbital. This type of bond is called a co-ordinate covalent bond. Although it is formed in a different manner than a regular covalent bond, once formed, it behaves in the same manner as any other single covalent bond.

needs lp

Question 29

Resonance Structures

Answer 29

One possible structure for the BF3 molecule is as follows:
( one bonds to f becomes double)

However, the central atom in this structure does not have a complete octet. If one pair of e- is moved from an F atom, making a double bond, then the B atom will have access to a complete octet.

In this structure, the B atom has a complete octet, even though the formal charges are not favourable. This structure would lead us to believe that there is 2 single B–F bonds and 1 double B=F. However, experimental evidence shows that the actual bond lengths within the BF3 molecule are all equal. This means that the bonds have properties that are somewhere between a single bond and a double bond.

This phenomenon is represented by drawing all the possible Lewis structures, and separating them by double-headed arrows.

Each of these structures is called a resonance structure. Notice that each structure shows the same relative position of the atoms in the molecule, but shows the different places of their bonding and lone pairs.

occurs when there is more than one way of positioning the DB –> resoance structures show all possibilities

Resonance occurs when we can draw two or more legitimate Lewis structures for the same molecule. Resonance only occurs when a molecule has at least one double bond. Molecules with only single bonds never show resonance. Resonance structures must all have the correct number of electrons and must all obey the octet rule.

bond length is the energy require to break atoms. pi bonds are not as strong but it does squish the bonds so the atrativness starts to increase because the lengths between bonds decreases

single bond: longest bond length, lowest bond energy and strong bond strength
DB: mid bond length, mid bond energy and stronger bond strength
TB: shortest bond length, highest bond energy and strongest bond strength

benzene is between single and double bond (1.5)

in benzene the db can inter convert between different bonds which is the db/e start to circle around

lp comes in and other comes out –> in the resonance hybrid you leave lp but draw dotted lines where the db would move and a negative 1/how many atoms its spread out on charge is added

charge is normally added to atom with most lp

Question 30

Bond and Molecular Polarity

Answer 30

Chemical bonds:
forces holding particles together in a compound. basis of chemical properties and reactions

Covalent Bonds:
Formed between non-metals. Electrons are shared by atoms. Driving force is the tendency to gain a stable octet.

Bond Polarity:
- Electronegativity – measure of the tendency of an atom to attract electrons towards itself when forming a covalent bond. (values on P.T.)
- Diatomic gases have 2 identical atoms involved in bonding and so their electronegativity values are exactly the same.
- Molecules with different types of atoms have different electronegativity values thus the sharing of electrons is not equal
- Polarity – in a covalent bond, the shared electrons are more the property of one atom than the other. Therefore one atom has a slightly more negative charge than the other atom which has a slightly more positive charge. These slight positive or negative charges are referred to as partial charges and are represented by the Greek letter delta.
- these partial charges are smaller than ionic compounds (partially ionic)
- Polarity determined by taking the difference of electronegativity values: Ionic bond = Greater than 1.7
Polar Covalent bond= Between 0.5 – 1.7
Non-Polar Covalent bond = Less than 0.5 to 0

MOLECULAR POLARITY:

• Lone pair: electron pairs not involved in bonding
• Bonding pair: electron pairs that are involved in bonding
• Electron arrangement around molecules so that they are a maximum distance from each other

Some molecules have a difference in electronegativities such that the atoms involved in the bond have partial charges.

DIPOLE MOMENTS: is used to predict the polarity of an entire molecule (not individual bonds). the dipole moment is a vector quantity (it has both direction and magnitude).

Given the shape (use VSEPR Theory) of a molecule, one can predict whether the molecule is polar or non-polar using the following steps:
- Determine whether or not there is a difference in electronegativities. If the answer is no, then the molecule is non-polar.
- If the answer is yes, then draw arrows such that the arrow head points towards the more electronegative atom (-|—>).
- Given the shape of the molecule and the direction of arrows: if the arrows point in opposite directions then there is a cancelling effect. If the arrows point in the same direction then there is an additive effect.
- if the arrows cancel each other in the molecule a non-polar molecule results. If the arrows add and point in a particular direction, then the molecule in question is a polar molecule.

Dipole moments can be used to distinguish between molecules that are made up of the same atoms but connected in a different way. Also polar molecules are affected by external electric fields.

Summary:
- A molecule that contains polar bonds will not necessarily be a polar molecule.
- If the shape of a molecule is such that the polar bonds are symmetrically charged, their effects many cancel and the molecule will be considered non-polar.
- Symmetrical molecules are non-polar.
- Symmetrical molecules include: Linear, Trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

what u do is draw the lewis shape, determine vsepr shape, add dipoles if en values show polar, examine dipoles and if adding then polar if cancelling then non-polar

no dipoles for non-polar or ionic compounds (based on en different)

in molecules where the atom is bonded to diff types of atoms you have to see the bond polarity of each type first –> even if dipoles cancel each other, if one bond is more polar based on en values then the cancelation is not equal and a bond may still have a dipole –> ex. cf2cl2 –> the c and f bond is very polar compared to the c and cl and when u draw dipoles it seems like non polar but in reality its polar because of difference

small dipoles may exist if the en difference is very close to being polar

bond polarity is polarity of bonds while molecular polarity is the polarity overall of a molecule

dipole moments used to predict the polarity of the overall molecule. show the biases in the bond. they are vector quantities so they tell us direction and magnitude/other info –> if they move in the same direction then additive affect (like upwards) but if they are in the opposite way and have the same energy/size then they cancel. shape of the molecule will also affect the bias. in diatomic there’s no bias so no neg or pos poles. Dipoles need to be equal in magnitude and direction for cancellation (which is why symmetry applies). When there diff atoms attached the en values may differ for the bonds so even when dipoles go opp direction they may not cancel due to the magnitudes not being the same.

what you do is you get the molecule find the bond. if non polar or ionic then u stop there. if polar then follow vsepr and add in dipoles. if cancel then non polar and if adding then polar.

there are no dipoles for lp.

with cis and trans isomers polarity can be different which is why electric fields can be used to determine polairty of each. bring electric field to see polarity. non polar do not react will polar will reflect the electric field.

if dipoles are equal in opp direction then they cancel. lone pairs are sometimes an indicator of polairty because symmetry becomes messed up but not all the time.

in the trigonal shape since the bond angles are all 120 degrees apart if the dipoles are all moving in opp direction and are equal they cancel.

EXCEPTION TO EN DIFFERENCES: bf3 is polar even though seems ionic. hf also polar even though seems ionic.

when drawing ion stuff don’t make a line just write the element symbol and + or - charges.

Question 31

extra vsper and lewis diagram note

Answer 31

when making molecule first count the total e- through adding the valences. then make rough sketch of e- to see if any lp (also can tell by doing total e- - highest multiple of 8 that underneath then divide by 2). then draw the correct lewis structure like linear if two attachment (not including lp) or other vsepr shape. make sure lp are added on equatorial for trigonal bipryamidal because of the 120 degree bond angle. then write the AXE notation. then write epg then molecular shape based on lp affect. for 3d structural diagram you basically make the wedge diagram except for trigonal planar or linear epg. if bent but tetrahedral then draw the two ligands far apart with bent shape and add lp with wedge at top left and right. make sure to show double bonds if they exist. show wedges on equatorial for trigonal bipyramidal and octahedral or just draw the traingle shape stuff.

Question 32

Intra particular forces

Answer 32

Intramolecular Forces: forces within a molecule that hold the atoms together in a chemical bond

Network Crystals (Covalent crystals):
- covalently bonded (intramolecular forces) atoms in a 3-D crystal
- very hard
- high melting point and high boiling points
- these crystals are the hardest to break since all bonds are covalent.
- diamond don’t conduct electricity but graphite does because the single lone electron
- all have repetitive structures/extensively covalently bonded structures
Examples of top 5 crystals/top in intra forces: Diamond C4, Graphite C3, Silicon dioxide SiO2, BN, SiC

Ionic Crystal (electrostatic attractions):
- 2nd stronger intraparticular force
- crystal lattice - cations and anions
- arranged to maximize attractive forces; regular repeating arrangement of ions
- anions and cations held together by electrostatic interactions
- strength of the ionic crystal is due to the fact that a lot of energy is required to pull away one ion from many.
- Lattice energy - the energy liberated from one mole of ionic crystal (one mole of solid ionic compound is completely separated into its gaseous ions). This is an experimentally determined quantity that shows the stability of the crystal.
- Hard
- High melting point and high boiling point
- strength of the ionic crystal depends on the charge of the cation, the higher the positive charge the more the attraction. Formula to determine strength: F = kq1q2 / r2. its basically the cation*anion over size. you basically do the absolute value of the charge of the cation * charge of anion. the one with larger charge indicates stronger bond but if same then you look at the size of the ions to see which is smaller as it will be stronger attraction.
- if hit by a hammer the ions would shift and then shatter. the hardness is because of electrostatic attractions and intraparticular force.
- Conclusions: When considering bond strength and crystal stability must take two factors into account: Charge and Radius
- Therefore, the larger the ionic radius, the weaker the force of attraction and the smaller the ionic radius, the stronger the force of attraction
- Note: when the charges of two ions being investigated are the same, then the radius of the ions involved becomes crucial in determining the degree of force of attraction.
- KEY: Bond strength and stability is proportional to the force of attraction. Thus the greater the force of attraction, the stronger the bond and the greater the force of attraction the higher the stability.
- when high bond strength then higher lattice energy and mp and bp but lower solubility

Metallic solids:
- Consist of atoms of metals held together via metallic bonding
- depends on the electronic structure of the metals.
- array of metal cations in a ‘sea’ of valence electrons
- example of metallic solids - zinc, iron
- example of mixtures of metals (Alloys): stainless steel, brass.
- two types of alloys: substitutional alloy (replacement) and interstitial alloys (small).
- the valence e- are mobile and are attracted by multiple cations/nuclei
- the higher the zeff the higher strength bond –> e- simultaneously being attracted by multiple nuclei. the higher the zeff the stronger the metallic bond. (Zeff = protons/atomic number - shelidibg electrons)

Question 33

Intermolecular forces

Answer 33

Intermolecular Forces: attractive forces between two distinct molecules

All solids, liquids, and gases are composed of atoms or molecules that must attract each other. When compounds are in a solid form the attraction between molecules is at its highest and the molecules are tightly packed together. But what forces govern these attractions?

Intermolecular bonding forces are present between atoms, molecules and ions results from simultaneous attraction of molecules.

Dipole-dipole forces - polar
London dispersion forces - non polar
Hydrogen bonding - extreme cases of dipole dipole (NOF)

Dipole – Dipole:
- This type of force is dominant in polar molecules (uneven dispersal of electrons)
- the presence of molecular dipoles causes simultaneous intermolecular attraction. The positive side of one molecule attracts the negative side of a neighbouring molecule.

Hydrogen bonding (weaker than covalent bonds):
- involves a special category of stronger dipole-dipole forces that exist between molecules containing Hydrogen atoms bonded to particularly electronegative elements
- hydrogen bonds are not really ‘hydrogen bonds’ – they are simply attractions between highly polar molecules.
- intermolecular forces that exist between polar molecules containing hydrogen atoms (slightly positive) and atoms such as O, N, and F (slightly negative)
- the hydrogen bond is weaker than the intramolecular bonds of the molecule.
- intense dipole between h and NOF so intense hand holding
- need lp on NOF for h bonding
- The solid line represents covalent bonds
- The dashed lines represent hydrogen bonds
- alcohols, ca, amines and amides of 1 and second degree exhibit H bonding

Properties of water:
- Exists as liquid, solid and gas naturally
- Most common substance on earth
- It is the universal solvent
- High specific heat due to hydrogen bonding (the ability to absorb a lot of heat with very little change in temperature since hydrogen bonds must be broken)
- Ice is less dense the liquid water due to hydrogen bonding (air pockets form which makes it less dense
- High boiling point and melting point due to hydrogen bonding

London dispersion forces (instantaneous dipoles):
- Occurs in non-polar molecules and monoatomic noble gases
- these are intermolecular forces that do not depend on polarity
- instantaneous uneven distribution of electrons allows the molecules to line up and attract neighbouring molecules
- all molecules polar and non-polar exhibit London dispersion forces
- Molecular or atomic crystal: considering melting point when in the frozen state. Example of a molecular crystal is CO2. Example of a atomic crystal is Ne, He etc
- based on e- distribution when uneven then LDF happens
- the higher the total number of electrons the stronger the intermolecular force/LDF

Molecular solids:
- Consists of molecules held together by intermolecular forces (hydrogen bonding in ice)
- Example: wax and ice

SUMMARY:
- network crystals: C4, C3, Sio2, BN, SiC
- ionic crystal: charge and then shape
- metallic solid: increase in zeff

• H- bonding: H+ NOF
• Dipole dipole
• LDF

usually ldf is stronger than dipole dipolse but DD is stronger if the same mass

to predict the intermolecular forces present see if NOF +H for hydrogen bonding, anything part of a molecule which would be polar for dipole dipole and if anythings non polar for LDF.

From highest boiling point to lowest: network crystal (c4,c3,sio2,sic, and bn), ionic crystal (look at charge and size (higher charge more stronger and smaller molecule more stronger), metallic solid (increase in Zeff or less shielding means higher bp since electrons are held close), hydrogen bonding (NOF+H but need lp), dipole dipole (polar molecules) and LDF (non polar or monoatomic = NG and if more than one non polar molecule then see one with more number of electrons, the one with more has a higher bp)

Question 34

Dissolving and extra facts about carbon containing molecules

Answer 34

Like disolves like so therefor polar disolves in polar and ionic compounds and non polar disolves in non polar. ionic only disolves when it is soluble (based on solubility rules). with ionic compounds dissolving there is ion dipoel attractions

we do not consider intermolecular forces in ionic bonding because there are no molecules –> we see that the type of intramolecular bond determines the type of intermolecular force. Oil and water do not mix as the partial charges on water attract pushing the oil with no partial charges out of the way.

Diamond versus Graphite:
- Diamond and graphite are both made up of carbon atoms that are bonded in different ways. This difference in arrangement is responsible for their unique properties. Thus diamond and graphite are allotropes.

• Allotropes - different forms of the same element.

Properties of diamond:
- All four valence electrons in carbon are involved in bonding.
- Each carbon atom is bonded to 4 other carbons.
- Hardness due to the strong covalent bonds.
- Does not conduct electricity (no electrons/lp available)

Properties of Graphite:
- Only 3 of the 4 valence electrons are involved in bonding.
- Carbons are arranged in six-membered rings to form layers. (two dimensional layers)
- The layers are held together by intermolecular forces.
- Covalent bonds account for hardness.
- It is slippery to the touch due to the fact that the layers slide past each other easily because of the weak intermolecular forces.
- Does conduct electricity (free electrons in structure)
- Insoluble in water and organic solvents.

Properties of Silicon Dioxide:
- Crystalline silicon has the same structure as diamond.
- Silicon also has 4 valence electrons.
- Silicon dioxide is a 3-D giant structure with each silicon bonded to oxygen atoms.
- High melting point due to strong covalent bonds that must be broken.
- Does not conduct electricity – there are no free electrons to move around.
- Insoluble in water and organic solvents (no attraction strong enough to overcome covalent bonds in giant structure).

Question 35

Extra notes on intermolecular forces

Answer 35

Types of Intermolecular Forces

Named for the Dutch scientist Johannes Direrik van der Waals, a physics professor at Amsterdam university, van der Waals forces are present between atoms, molecules and ions. Van der Waals forces are subdivided into dipole-dipole forces and London dispersion forces.

1. Dipole -Dipole Forces:
   - Found in a molecule with a charge separation (a dipole)
   - If the molecules in a sample are polar (uneven dispersal of electrons) the presence of molecular dipoles causes simultaneous intermolecular attraction. The positive side of one molecule attracts the negative side of a neighboring molecule. Molecules line up in crystal such that opposite dipoles are next to each other.
   - Result is molecules are held together and an increased amount of energy is required to separate molecules (i.e. higher b.p)
   - Strength of attraction depends on the size of the charges on the molecules.
2. London Dispersion Forces (a.k.a dispersion forces)
   - In non-polar molecules and monoatomic noble gases these are intermolecular forces that do not depend on polarity. Nuclei attract more than just the electrons that form the bonds of the molecule; electrons from neighboring complexes also become attracted. A temporarily induced dipole because of the molecule moving. (uneven distribution)
   - As molecules moves, electron density shifts (like water in a balloon) creating a dipole. At that point small dipole-dipole interaction occurs.
   - Two factors influence the strength of LDFs: size (# of electrons, more the stronger ldf) and shape (affects cozy up ability)
   - The number of electrons in the molecules - the more electrons the stronger the London dispersions forces.
   - The shapes of the molecules - shape affects how closely the molecules may approach each other in solid and liquid states.

Note: London dispersion forces are weaker than covalent bonds. This is because the simultaneous attracted electrons further from the nuclei than electrons in covalent bonds and there is also repulsion due to the shielding effects from inner electrons. All molecules have London dispersion forces, and if they are polar they will also have dipole-dipole forces also. Of the two London dispersion forces are more significant.

3. Hydrogen Bonding (H + N,O,F)
   - Hydrogen bonding involves a special category of stronger dipole-dipole forces that exist between molecules containing Hydrogen atoms bonded to particularly electronegative elements. bonds are not really ‘hydrogen bonds’ - they are simply attractions between highly polar molecules.
   - Do not involve sharing of electrons (not covalent) nor does it bind two ions together (not ionic)
   - Is a special case of dipole-dipole interaction but much stronger due to the high electronegativity associated with Fluorine, Oxygen and Nitrogen
   - To have hydrogen bonding you MUST have a lone pair of electrons.

All solids, liquids, and gases are composed of atoms or molecules that must attract each other. When compounds are in a solid form the attraction between molecules is at its highest and the molecules are tightly packed together. But what forces govern these attractions? Solids (crystals or amorphous) are the result of ionic bonds, covalent bonds, intermolecular attractions between molecules or metallic bonding

Structures of Solids/Crystals: the arrangement of ions/atoms/molecules can result in 2 possible shapes:
1. Crystalline solid:
- ordered in well-defined arrangements of particles. (flat surfaces or regular shapes)
- examples: diamond, quartz, salt
2. Amorphous solid:
- Irregular or random arrangement of particles (no orderly structure).
- Examples: glass, rubber

Like covalent bonds, intermolecular bonding results from simultaneous attraction of molecules. however, there is an important difference between intermolecular bonding and covalent bonding.
- Valence electrons are not shared – bonding capacity is complete in forming molecules.
- intermolecular forces attract as many molecules as can physically surround one another – three-dimensional.
- Intermolecular forces are much weaker than covalent(intra) bonds.

Effect on Physical Properties:
Since melting and boiling involves the separation of molecules from one another, it makes sense that any substance with intermolecular forces would have an increased melting and boiling point.

Melting point: solid to liquid
Boiling point: liquid to gas

Molecular liquids:
- Volatity (volatile): ablity of a liquid to vaporize. In volatile liquids, the molecules have enough energy to escape the liquid phase and enter gaseous phase. Vaporizing molecular substances need to overcome intermolecular forces. For volatie liquids, their intermolecular forces are weak and require a low amount of energy to overcome intermolecular forces. Larger molecules will have lower volatility as compared to smaller molecules.

Molecular crystals:
- Properties of molecular crystals (covalaently bonded crystals)
1. Neither liquid nor solid conduct electricity. (besides graphite and bucky ball)
2. Many exist as a gas at room temperature and atmosphere.
a. Those that are solid or liquid are relatively volatile.
3. Melting point of solid crystals are relatively low.
4. Boiling point of liquids are relatively low.

5. Solids are generally soft and have a waxy consistency.
6. Large amount of energy is required to chemically decompose a substance into simpler substances. (arm ripliy)

Notes:
- Point #1 refers to both polar and non-polar molecules –> Polar molecular crystal: molecules on the lattice points are polar. Non-polar molecular crystal: molecules on the lattice points are non-polar
- Points #2-5 refer to intermolecular forces (which is why there is low bp and mp)
- #6 refers to intramolecular forces

Intramolecular Forces: Force WITHIN an individual molecule
* covalent, network crystal, ionic crystal and metallic solid
Intermolecular Forces: Forces BETWEEN molecules –> h bonding, d-d, and ldf

nonpolar: larger molecular size, more electrons, greater polarizability (instantaneous dipoles), stronger ldfs, more energy needed to overcome IMF’s so higher bp and mp

polar: less strong that h bonding. dipole dipole attractions require energy to break. the higher the en difference, the larger dipoles and stronger attractions so more energy needed to break.

h bonding: extreme dipoles that results in strong intermolecular attraction lots of energy needed to break. higher en difference more extreme dipole so stronger attractions more energy required to break.

metallic solid: held together by positive cations attracting mobile valence e-, energy required to separate atoms. more zeff means more effective nuclear nuclear charge and attraction so the valence e- are more attracted by the positive cation. lower shielding means that the valence e- are held close so stronger attraction and more energy required to pull separate atoms.

ionic crystal: held together by strong electrostatic attractions so high energy is needed to separate ions. the higher the force of attraction (based on charges) the more energy needed to separate ions since bond strength is proportional to force of attraction. the smaller the size of the cation and anion the stronger the fore of attraction since the ions can get close to each other and cozy up.

pentane and enopentane have the same chemical formula, so they are structural isomers, have same number of e-, non polar and exhibit ldfs. but they have different mp and bp because pentane has a straight chain which allows for greater surface area available for ldf so higher bp and neopentane has a more compact structure allowing for better packing so higher mp

Question 36

explanations for bp and mp order (highest to lowest)

Answer 36

Any network crystal: 3-d covalently bonded crystal made up of carbon, strength due to strong covalent bonds (intra) and therefore high energy required to boil

ionic crystal: ionic crystal held together by strong electrostatic attractions, high energy needed to separate ions but not as much as network crystals

metallic solid: metallic solid held together by positive cations attracting mobile valence e-, energy required to separate atoms which is high but not as much as ionic crystals

polar molecules with h bonding: polar molecule with an extreme dipole that results in strong intermolecular attraction called hydrogen bonding, lots of energy required to separate or break hydrogen bonds but not as strong as an intraparticular attraction

polar only dipole dipole: polar molecule without an extreme dipole, dipole dipole attractions require less energy to break intermolecular attractions as compared to hydrogen bonding

non polar molecule, ldfs: non polar molecule, exhibits ldf’s due to the instantaneous dipoles between molecules, weak intermolecular force therefore less energy required to boil or melt

when two of the same type give extra examples as to why they are placed before the other example in two non polar molecules the one with more e- is placed higher as with more e- there is greater polarizability and stronger ldfs and therefore more energy needed to overcome the IMF’s. or for metallic solids you would look at zeff, higher zeff means the electrons are held more closer so harder to pull apart or if there’s more less sheilding its more closer to nucleus so attractions are more stronger. for ionic cyrtslas look at charges and size.

when multiple non polar look at TOTAL number of e-, if multiple metallic and zeff is also the same then look at shielding because when shielding is low there is stronger metallic bond because stronger attraction between the positive cation and the mobile valence e-. when looking at multiple ionic if the charge is same then look at size of the ions, smaller cation and anion results in a higher bp or mp because can hold together more close.